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REESE   LIBRARY 


UNIVERSITY  OF  CALIFORNIA. 
MAR  11  1893 

CLm  No.... 


From  the  collection  of  the 

7   n 

P   T   m 
0  Jrrelinger 

v    JLJibrary 


San  Francisco,  California 
2006 


LABORATORY  MANUAL 


OF 


CHEMISTRY 


BY 

JAMES   E.  ARMSTRONG 

PRINCIPAL    OF    ENGLEWOOD   HIGH    SCHOOL,    CHICAGO 

AND 

JAMES  H.  NORTON 

PRINCIPAL   OF    LAKE   VIEW    HIGH    SCHOOL,    CHICAGO 


NEW  YORK  •:•  CINCINNATI  •:•  CHICAGO 

AMERICAN     BOOK     COMPANY 


COPYRIGHT,  1891,  BY 
AMERICAN   BOOK  COMPANY 


[All  rights  rcserved\ 


Printed  by  Wm.  Ivisoa 
New  York,  U.S.A. 


PREFACE. 

"TN  preparing  this  brief  course,  the  authors  have  endeavored  to  select 
-•-  such  work  as  will  best  develop  the  true  spirit  of  investigation.  It 
is  not  intended  to  present  an  exhaustive  treatise,  but  rather  a  few  of 
the  stepping-stones  over  which  the  student  of  this  department  of 
nature  must  travel.  We  most  thoroughly  believe  that  the  experiments 
selected  should,  as  a  rule,  be  of  such  a  nature  that  the  pupil  can  perform 
them  for  himself.  The  teacher  should  be  but  the  guide  that  points 
out  the  right  path,  calling  attention  to  the  by-paths  of  error. 

The  work  here  laid  out  should  be  accompanied  by  some  text-book. 
The  authors  have  had  in  mind,  while  preparing  this  work,  the  "  Manual 
of  Chemistry,"  by  Eliot  and  Stofer ;-  although  it  could  be  used  with  any 
other  good  text-book  of  elementary  chemistry.  It  is  our  plan  to  require 
three  hours'  work  per  week  in  the  laboratory  and  two  in  the  recitation- 
room.  The  experiments  here  given,  with  the  necessary  practice  in 
determining  unknown  substances  by  the  Key,  will  probably  occupy  the 
class  for  forty  weeks.  Some  experiments  should  be  indicated  by  the 
teacher  for  omission,  if  the  time  alloted  for  chemistry  is  too  short. 

We  believe  the  best  results  will  be  attained  by  requiring  the  pupil 
to  write  his  notes  in  the  laboratory  at  the  time  the  experiment  is  per- 
formed, and  not  after  he  has  been  assisted  to  draw  "  bookish  "  conclu- 
sions from  various  sources.  If  the  work  is  not  so  neatly  done,  or  if 
wrong  conclusions  are  reached,  the  mistakes  are  his  own  and  can  be 
corrected  by  pointing  out,  not  the  mistaken  inference  but  the  errors 
in  conditions,  accidents  and  coincidences.  Let  the  pupil  perform  the 
experiment  again  in  the  teacher's  presence,  or  perform  it  by  direction 
under  new  conditions,  and  so  finally  be  led  to  detect  the  error  of 

judgment.     Long  descriptions  of  experiments  are  usually  unprofitable  to 

iii 


IV  PREFACE. 

both  pupil  and  teacher.  A  simple  statement  of  the  results  obtained, 
containing  the  answers  to  the  questions  asked,  is  usually  all  that 
should  be  given,  unless  otherwise  directed.  When  the  teacher  performs 
an  experiment,  a  full  account  of  the  process  should  be  written.  The 
"Laboratory  Manual"  should  not  be  taken  from  the  laboratory,  as  it 
offers  opportunity  to  copy  work  that  should  be  done  experimentally. 

We  wish  to  acknowledge  our  obligations  to  Professor  A.  V.  E. 
Young,  Northwestern  University,  Evanston,  111.,  and  to  Professor 
Le  Roy  C.  Cooley,  Vassar  College,  for  special  experiments  noted  in  the 
text ;  to  Professor  F.  Sanford,  Leland  Stanford  University,  California, 
Professor  M.  Delafontain,  Mr.  F.  L.  Morse  and  Mr.  A.  L.  Smith,  of  the 
Chicago  High  Schools,  for  valuable  suggestions  in  preparing  the  work;  and 
also  to  Mr.  W.  E.  Danner,  of  the  Nutriment  Company,  Union  Stock 
Yards,  Chicago,  for  materials  and  suggestions  in  preparing  the  experi- 
ments upon  Digestive  Ferments. 

CHICAGO,  ILL.,  October,  1891. 


LIST   OF   EXPERIMENTS. 


EXP 
1. 


27. 


Physical  Properties  of  Matter 

Specific  Volume 

Specific  Gravity 

Changes  due  to  Heat 12 

Latent  Heat  

Changes  affecting  the  Boiling  Point 

Evaporation,   Vaporization,    Distil- 
lation  

Chemical  Change,  Sulphur  and  Iron. 

Chemical    Change,     Electrolysis   of 
Water 

Solution 18 

Effect  of  Heat  upon  Solution 20 

Crystallization 20 

Latent  Heat  of  Solution,  Ice  and  Salt 

Latent     Heat   of  Solution,    Sodium 
Phosphate  and  Nitric  Acid 

Allotropism 

Effect   of    Solution  upon  Chemical 
Change 

Divisibility  of  Matter,  Molecules 

Porosity  of  Matter,   Proximity   of 
Molecules 

Atoms 

Compounds  and  Elements 

Compounds  from  so  few  Elements, 
How  possible 

Analysis 28 

Synthesis 

Acids  and  Alkalies 

Neutralization,    Formation    of 
Salts 

Indestructibility    of    Matter,     Elec- 
trolysis of  Water 36 

Indestructibility  of  Matter,  Mercuric 
Sulphocyanate  . 


GE 

EXP. 

8 

28. 

8 

29. 

10 

30. 

12 

31. 

12 

32. 

14 

33, 

34. 

14 

16 

35. 

18 

36. 

18 

37. 

20 

38. 

20 

33. 

20 

40. 

22 

41.  ] 

22 

42.  ] 

43. 

24 

44.  ] 

24 

45.  ] 

43.  ] 

24 

47.  . 

26 

48.  1 

26 

49.  ( 

50.  1 

28 

51.  1 

28 

52.  ] 

30 

53.  ( 

30 

54.  1 

55.  1 

32 

56.  1 

57.  1 

36 

58.  ( 

36 

59.  i 

Law  of  Definite  Proportions 
Law  of  Multiple  Proportions 
Chemistry  of  Water,  Composition.  . 
Distillation  of  Water 
Water  of  Crysta'lizaf  ion 
Efflorescence  and  Deliquescence 


PAGE 

3d 
38 

..  40 
40 
42 
42 


Impurities    of  Drinking-water,     Or- 

ganic ....................  ..........  42 

Impurities  of  Drinking-water, 

Chlorine  .........................  44 

Hydrogen  ...........................  44 

Hydrogen,  continued  ................  46 

Hydrogen,  continued  ................  46 

Synthesis  of  Water  by  Volume, 

Valency  .........................  48 

Chemistry  of  Air,  Weight  of  Air.  ....  50 

Heating  Lead  Oxide  .................  50 

Preventing  Oxidation  ................  50 

Nitrogen  of  Air  ................  "  .....  52 

Nitrogen  ............................  52 

Nitrogen,  continued  ...............  .  .  52 

Per  cent,  of  nitrogen  in  air  .........  52 

Analysis  of  Air  ......................  54 

Moisture  in  Air  ......................  58 

Carbon  Dioxide  in  Air  ..............  58 

Preparation  of  Oxygen  ..............  60 

Properties  of  Oxygen  ...............  60 

Effect  of  Oxygen  upon  Blood  .......  60 

Oxygen  from  Green  Plants  ..........  62 

Nitrous  Oxide  .......................  62 

Nitrous  Oxide,  continued  ............  62 

Nitric  Oxide  .........................  64 

Nitric  Oxide,  continued  ..............  6  1 

Compound  of  Nitrogen  and  Hydro- 

gen Ammonia  .....................  66 

Ammonia  from  Ammonium  Chloride  6C 


LIST  OF  EXPERIMENTS. 


EXP.  PAGE 

60.  Chlorine 66 

61.  Chlorine,  continued 68 

62.  Chlorine  Water  08 

63.  Bromine .  70 

64^,  Iodine 70 

65.  Hydrochloric  Acid 70 

66.  Characteristic  Test  for  Hydrochloric 

Acid 72 

67.  Nitric  Acid  and  Characteristic  Test. .  72 

68.  Sulphuric    Acid  and  Characteristic 

Test 72 

69.  Hydrosulphuric  Acid  and  Character- 

istic Test 74 

70.  Carbonic   Acid   and    Characteristic 

Test 74 

71.  Sulphur  Dioxide 76 

72.  Phosphorus 76 

73.  Arsenic 78 

74.  Test  for  Arsenic 78 

75.  Carbon  Dioxide        78 

76.  Preparation  of  CO2 80 

77.  Absorption  of  CO, 80 

78.  Saturation  with  CO, 82 

79.  Candle  Flame 82 

80.  The  Flame,  continued 84 

81.  Kindling  Temperature 84 

82.  The  Bunsen  Flame 84 

83.  The  Principle  of  Sir  Humphry  Davy's 

Safety-lamp 86 

84.  The  Action  of   Hot     Carbon   upon 

Oxygen 83 

85.  The  Blow-pipe 83 

86.  The  Oxidizing  Flame 88 

87.  The  Borax  Bead 88 

88.  Marsh  Gas 88 

89.  Illuminating  Gas 90 

90.  Alcoholic  Fermentation 90 

91.  Acetic  Fermentation 92 

92.  Fractional  Distillation C2 

93.  Test  for  Sugar £2 

94.  Digestive  Ferments,  Albuminous 94 

95.  Effect    of    Temperature    upon   Di- 

gestion    94 

96.  Effect  of  Mastication  upon  Diges- 

tion   94 

97.  Digestion  of  Starches  . .  90 

98.  Effect   of  Acids  and  Alkalies  upon 

Starch  Digestion 98 

99.  Digestion  of  Milk 98 

100.  Sulphides 100 

101.  Chlorides 102 

102.  Carbonates 102 


EXP.  PAGE 

103.  Blow-pipe  Tests,  Metals 104 

104.  Blow-pipe  Tests,  Salts 104 

105.  Borax-bead  Tests 104 

100.  Flame  Tests 104 

107.  Separation  and  Identification  of  First 

Group  Metals 108 

108.  Separation   and   Identification,  con- 

tinued    103 

109.  Separation  and  Identification,    con- 

tinued    108 

110.  Separation  and  Identification,    con- 

tinued   108 

111.  Confirmatory  Tests  of  First  Group...  110 

112.  Reduction  of  Silver 110 

113.  Pure  Silver  from  Silver  Coin 110 

114.  Silver-printing 112 

115.  Silver-plating 112 

116.  Reduction  of  Lead  by  Zinc 112 

117.  Distinction    between   Mercuric  and 

Mercurous  Salts 114 

118.  Reduction  of  Mercury 114 

119.  Separation   of    Second   Group    into 

two  Divisions 114 

120.  Separation    of    Second    Group   into 

two  Divisions,  continued HO 

121.  Separation    of    Second    Group   into 

two  Divisions,  continued 116 

122.  Separation    of    Second    Group   into 

two  Divisions,  continued no 

123.  Identification  of  Antimony lie 

124.  Identification  of  Tin 110 

K5.  Identification  of  Tin,  continued 116 

126.  Identification  of  Copper 118 

127.  Confirmatory  Test  for  Copper 118 

128.  Bismuth 118 

129.  Separation     and     Identification    of 

Metals    of  Second  Division   into 
Second  Group 120 

130.  Iron 120 

131.  Iron,  continued 120 

132.  Iron,  continued 122 

133.  Iron,  continued 122 

134.  Iron,  continued 122 

135.  Iron,  continued 124 

136.  Aluminum 124 

137.  Chromium  ...     124 

138.  Separation  of  Metals  of   the   Third 

Group 120 

139.  Separation   of   Metals  of  the  Third 

Group,  continued 120 

140.  Separation  and  Identification  of  Iron, 

Aluminum,  and  Chromium 126 


LIST  OF  EXPERIMENTS. 


EXP.  PAGE 

141.  Separation  and  Identification  of  Iron, 

Aluminum,  and    Chromium,    con- 
tinued    126 

142.  Confirmatory  Test  for  Chromium.   .   120 

143.  Confirmatory  Test    for   Chromium, 

continued 126 

144.  Confirmatory  Trst  for  Aluminum 12»i 

145.  Separation  and  Identification  of  Ni, 

Co,  Mn,  and  Zn 128 

146.  Separation  and  Identification  of  Wi, 

Co,  Mn,  and  Zn,  continued 128 

147.  Separation  and  Identification  of  Wi, 

Co,  Mn,  and  Zn,  continued 128 

148.  Separation  and  Identification  of  Ni, 

Co,  Mn,  and  Zn,  continued 128 

149.  Separation  and  Identificalion  of  N"i, 

Co,  Mn,  and  Zn,  continued     ...  128 

150.  Separation  and  Identification  of  Ni, 

Co,  Mn,  and  Zn,  continued 128 


EXP  PAGE 

151.  Fourth  Group  Metals,  Calcium  Oxide  130 

152.  Fourth      Group     Metals,      Calcium 

Hydrate 130 

153.  Fourth  Group  Metals,  Calcium  Sul- 

phate   130 

154.  Fourth  Group  Metals,  Calcium  Flame 

Test 130 

155.  Fourth  Group  Metals,  Calcium  Car- 

bonate     130 

156.  Other  Carbonates  of  Fourth  Group..   132 

157.  Other  Carbonates  of  Fourth  Group, 

continued  1 32 

158.  Separation  of  Ba,  Sr,  and  Ca 132 

159.  Magnesium 13 -1 

160.  Ammonia 134 

161.  Sodium 134 

162.  Potassium 134 

163.  Potassium,  continued  134 

164.  Potassium  from  Wood  Ashes 134 


INTRODUCTORY  WORK. 


PHYSICAL    PEOPBRTIBS    OF    MATTER. 

Experiment  1. 

Materials  and  Apparatus. — Pieces  of  wood,  iron,  lead,  wax,  putty, 
glass,  sulphur,  .and  sugar. 

Examine  each  of  the  above,  and  compare  with  water  and  air. 

Compare  with  reference  to  color,  odor,  hardness,  brittleness,  elasticity, 
and  form. 

Do  they  occupy  space  1 

Have  they  weight  ? 

Do  they  appeal  to  each  of  the  senses  ?     Define  matter 

Experiment  2. —Specific  Volume. 

Materials  and  Apparatus. — Same  as  in  Experiment  1  and  metric  rule, 

balances,  pocket- 
knife,  test-tube, 
and  burette. 

(a)  Allow  5 
or  10  c.  c.  of  water 
to  run  into  a  test- 
tub  e  from  the 
burette.  Note  ex- 
act point  at  which 
water  stands  in 
your  tube. 

Pour  back  the 
contents  of  your 
test-tube  into  the 

Fig.  I.-CHEMICAL  BALANCE,  burette. 

Does  the  water  in  the  burette  stand  where  it  did  at  the  beginning 
of  the  experiment  1  Explain. 


10  LABORATORY    MANUAL. 

Carefully  mark  off  on  the  side  of  one  of  your  test-tubes  the  point  to 
which  5  c.  c.  water  rises.  Also,  if  convenient,  mark  2.5  c.  c.  and  10  c.  c. 

N".  B. — In  all  subsequent  experiments,  when  you  are  requested  to  take  5  c.  c.  or 
10  c.  c.  of  any  liquid,  measure  out  the  amount  as  accurately  as  possible,  using  your 
graduated  test-tube. 

Measure  out  25  c.  c.  of  water  in  a  beaker  by  filling  your  graduated 
test-tube  five  times  to  the  5  c.  c.  mark.  Now  pour  this  water  into  the 
burette  and  note  how  much  you  must  allow  for  errors  in  your  work 
when  you  measure  out  large  quantities  of  water  in  this  manner. 

(b)  Make  cubes  of  wax,  putty,  lead,  and  wood,  each  weighing  a  gram. 
Measure  carefully  with  metric  rule  and  compare  volumes,  using  c.  c. 
for  the  unit. 

What  is  the  volume  of  1  gram  of  water  at  the  temperature  of  its 
greatest  density  1 

Weigh  5  c.  c.  of  water  in  test-tube ;  empty  and  weigh  tube.  What 
result  do  you  obtain  as  an  answer  to  the  preceding  question  ? 

Now  compare  volumes  of  above  substances  to  that  of  1  gram  of  water, 
and  state  their  specific  volume. 

Experiment  3. — Specific  Gravity. 

Materials  and  Apparatus. — Balances,  test-tube  (rather  large),  sulphur, 
or  some  other  substance  of  which  to  determine  the  specific  gravity. 

Find  specific  gravity  of  the  substance  given  you,  in  any  way  you 
choose. 

One  method. — Carefully  weigh  one  of  the  cubes  (b,  Exp.  2).  Fill  a 
beaker  so  full  of  water  that  nothing  can  be  added  without  causing  an 
overflow,  and  place  the  beaker  in  a  porcelain  dish. 

Now  immerse  the  cube  in  the  water  in  the  beaker  and  collect  the  over- 
flow in  the  porcelain  dish  and  weigh. 

What  was  the  exact  weight  of  the  cube  ? 

What  is  the  volume  of  the  water  which  overflows  1 

What  is  the  weight  of  this  volume  of  water? 

How  many  times  the  weight  of  the  water  is  the  weight  of  the  solid  1 
This  is  the  "  specific  gravity  "  of  the  solid. 

Another  method. — Weigh  10  or  15  grams  of  dry  sulphur  in  lumps. 
Weigh  test-tube  filled  with  water  and  corked.  Put  in  sulphur,  fill  with 
water,  and  weigh  again. 

Specific  gravity  =  weight  of  sulphur  -f-  weight  of  equal  volume  of 
water. 


12  LABORATORY    MANUAL. 

CHANGE. — Performing  an  experiment  in  physics  or  chemistry  consists 
in  surrounding  some  form  of  matter  with  certain  conditions,  which  should 
be  carefully  noted,  and  in  watching  the  result.  Accounting  for  the  changes 
which  have  taken  place  conies  next,  and  is  often  the  most  difficult 
portion  of  the  work. 

Experiment  4. — Some  Changes  Due  to  Heat. 

(a)  Fill  test-tube  with  water  so  that  it  stands  up  above  the  edge  of 
the  tube.     Warm  gently.     What  happens  ?     How  do  you  account  for  it1? 
What  change  takes  place  in  the  water?     Does  this  change   involve  a 
change  in  the  nature  of  the  substance  ? 

(b)  Fit  a  short  piece  of  glass  tubing  into  a  cork ;  till  test-tube  with 
water  and  insert  cork,  forcing  the  liquid  up  to  a  point  where  it  can 
be  seen  in  the  tube.      Warm  the  water  in  the  test-tube  as  in  (a),  and 
answer  the  same  questions. 

(c)  Perform  the  same  experiment,  having  air  in  the  test-tube  and  a 
drop  of  water  in  the  glass  tubing.    Answer  same  questions  as  in  (a).    This 
is  an  air  thermometer.     What  is  a  differential  thermometer  1 

Experiment  5. — Latent  Heat. 

Apparatus. — Ring-stand,  evaporating  dish  or  glass  beaker,  and  chemi- 
cal thermometer.  Take  a  quantity  of  pounded  ice  or  snow — enough  to 
fill  your  beaker.  Place  on  wire-gauze  or  sand-bath,  on  your  ring-stand, 
and  heat. 

Note  temperature  when  melting  begins. 

Note  temperature  while  melting  proceeds. 

Note  temperature  just  as  last  particle  is  melted. 

How  much  time  was  required  1     State  as  exactly  as  possible. 

Continue  to  heat  to  boiling  point.  Note  temperature  at  which  it 
begins  to  boil. 

Note  time  required  to  heat  the  water — after  the  ice  was  melted — to 
the  point  where  it  began  to  boil. 

Boil  water  all  away  and  note  time  required. 

Compare  the  three  results  obtained. 

How  do  your  results  differ  from  figures  given  in  the  text  ? 

What  is  latent  heat  ? 

State  all  the  changes  which  take  place  in  above  experiment. 

Is  the  substance  changed  ?     To  what  extent  ? 

Are  ice,  water,  and  steam  one  substance  ? 


14  LABORATORY   MANUAL. 


Experiment  6.  —  Changes  Affecting  the  Boiling  Point. 

(a)  Find  temperature  of  boiling  for  some  other  liquid. 

(b)  Make  strong  solution  of  salt  in  water  and  note  how  boiling  point 
is  changed. 

N.  B.  —Care  should  be  taken  in  the  following  to  avoid  explosion.  Note  these 
points  of  caution  : 

Just  as  soon  as  boiling  begins,  remove  heat  and  put  in  stopper.  In  inverting 
tube  be  careful  not  to  burn  your  hands.  Now  pour  on  a  quantity  of  cold  water, 
and  as  soon  as  water  in  tube  begins  to  boil  for  the  first  time,  so  that  you  are  able  to 
see  the  phenomenon,  bring  your  experiment  to  an  end  and  do  not  try  to  make  the 
water  boil  a  second  time. 

(c)  Boil  water  in  test-tube  (half  full),  insert  stopper,  and  invert. 
Remove  the  heat  and  bring  the  water  to  boil  again  by  pouring  on 

cold  water. 

How  does  the  cold  water  affect  the  matter  above  the  water  ? 
State  your  conclusion  in  regard  to  (u). 


Experiment  7.  —  Evaporation,  Vaporization,  Distillation. 

Boil  water  and  collect  steam  on  piece  of  cold  glass  or  porcelain. 

What  changes  take  place  ? 

What  is*  distillation  1 

Place  a  few  crystals  of  iodine  in  test-tube,  heat,  and  describe  changes 
which  take  place. 

Examine  substance  formed  on  the  sides  of  the  tube  near  the  top, 
where  it  is  cool. 

Does  it  resemble  the  substance  with  which  you  started  1 

What  is  sublimation  1 

To  what  extent  have  the  substances  been  changed  in  this  experiment  1 

QUESTIONS  ON  EXPERIMENTS  4,  5,   6,  AND  7. 

In  any  of  the  changes  produced,  has  the  nature  of  the  substance 
been  altered  1 

Put  a  few  drops  of  ether  on  your  wrist  and  blow  on  it  ;  account  for 
the  sensation  produced. 

What  does  this  suggest  with  regard  to  perspiration  1 

Does  melting  take  place  through  a  wide  range  of  temperature  ? 

Does  vaporization  1 

Which  form  1 


16 


LABORATORY   MANUAL. 


Experiment  8. — Chemical  Change. 

(a)  Try   to   dissolve  a  small   quantity  of  flowers  of  sulphur  in  bi- 
sulphide of  carbon. 

Try  to  pick  up  small  particles  of  iron  with  a  magnet,  using  iron  tilings 
or  powdered  iron.  This  shows  one  of  the  properties  of  iron  by  which 
you  can  recognize  it  in  the  future. 

Try  some  dilute  hydrochloric  acid  on  some  sulphur.  Is  any  gas 
given  off  having  an  odor  1 

Try  same  acid  on  particles  of  iron  and  answer  same  question  as  to 
gas  in  this  case. 

(b)  Mix  thoroughly  small  quantities  of   powdered  iron  and  flowers 
of  sulphur. 

Is  a  new  substance  formed  ? 
Is  the  sulphur  still  sulphur  1 
Try  bisulphide  of  carbon. 
Is  the  iron  still  iron  1 
Try  magnet. 

(c)  Weigh  out  3.2  grams  sulphur  flowers,  also  5.6  grams  powdered 

iron.  Mix  thoroughly  and  place  in 
hard  glass  test-tube  and  heat  over 
burner  until  the  mass  suddenly 
glows  vividly. 

Set  aside  to  cool.  When  it  has 
cooled,  examine  for  sulphur  and 
for  iron. 

Is  the  iron  still  iron  ? 

Is  the  sulphur  still  sulphur  '[ 

Try  substance  with  hydro- 
chloric acid  and  note  odor  of  gas 
given  off  and  its  effect  upon  paper 
moistened  in  solution  of  lead 
acetate. 

How  many  substances  had  you 
to  begin  with  1 
How  many  have  you  now  1 

How  does  this  change  differ  from  those  changes  noted  in  all  previous 
experiments  1 

With  what  kind  of  changes  is  chemistry  concerned  ? 


Fig.  2. — UNION  BY  HEAT. 


NOTE. --Whenever  an  experiment  is  marked  "For  the  Teacher,"  it  should  be 
performed  by  the  teacher  in  the  presence  of  the  class. 


18 


LABORATORY   MANUAL. 


Experiment  9. — Chemical  Change,  Continued. 

(For  the  Teacher.} — With  the  apparatus  shown  in  Fig.  3,  the  teacher 
is  to  perform  the  familiar  experiment  of  electrolysis  of  water,  while  the 
pupils  take  notes  of  all  that  goes  on  and  answer  on  opposite  page  any 
questions  which  the  teacher  may  emphasize. 

Note  following  points : 
How  many  substances  have  you  as  a 
result  of  your  experiment  ?  How  many  had 
you  to  begin  with  1  How  does  experiment 
in  this  respect  differ  from  preceding? 
What  is  analysis  1  What  is  synthesis  1 


Fig.  3. — ELECTROLYSIS  OF  WATER. 

NOTE. — The  above  experiment  is  here  introduced  merely  to  show  the  difference 
between  chemical  and  physical  changes.  The  analysis  of  water  will  be  discussed 
later. 

Experiment  10. — Solution. 

Materials  and  Apparatus. — Large  beakers,  sugar,  filter-paper,  funnel, 
funnel-stand. 

Take  50  c.  c.  of  water  in  a  beaker  and  add  to  it  10  grams  of  finely 
pulverized  sugar,  stirring  with  a  glass  rod.  Does  the  sugar  disappear  ? 
Where  has  it  gone  ?  Is  this  a  physical  or  a  chemical  change  ?  Has  the 
sugar  lost  its  essential  characteristics  1  Taste  it.  Fold  a  filter-paper 
and  place  it  in  the  glass  funnel.  Support  this  on  the  funnel-stand  and 
pour  the  solution  into  it.  Does  the  sugar  pass  through  as  readily  as 
water,  or  can  you  separate  out  the  sugar  1  Can  you  separate  them  by 
boiling  away  the  water?  Evaporate  until  you  get  a  thick  sirup,  and 
then  set  aside  for  several  days.  Can  you  dissolve  clay  in  water  ?  Try  it. 
Pass  it  several  times  through  two  thicknesses  of  filter-paper.  Is  this  a 
solution  1  What  is  a  solution  ? 


20 


LABORATORY    MANUAL. 


Experiment  11. — Hot  Solution. 

Materials  and  Apparatus. — Beaker,  alum,  mortar,  and  pestle. 

Take  25  grams  of  cold  water 
in  a  beaker  and  add  to  it  10  grams 
of  pulverized  alum.  Stir  for  some 
time.  Does  all  the  alum  dissolve? 
Why  will  it  not  dissolve  any 
amount  you  may  add?  A  sponge 
full  of  water  is  said  to  be  satu- 
rated, so  this  is  called  a  saturated 
solution.  !Now  heat  the  solution 
by  placing  the  beaker  upon  the 
sand-bath  over  the  gas-burner. 
Does  the  alum  all  disappear?  What 
do  you  infer  as  to  the  effect  of 
temperature  upon  the  amount  of 
alum  held  in  solution  ?  Save  the 
alum  solution  24  hours  or  more 

for  another  experiment. 

• 

Fig.  4.— FILTRATION. 

Experiment  12.  —Crystallization. 

Materials. — Sugar  solution  and  alum  solution  of  last  experiments. 

Examine  the  two  solutions  set  aside.  What  has  occurred  1  Has 
any  of  the  water  evaporated  ?  Examine  some  of  the  most  regular 
crystals  of  each.  Have  they  definite  forms  ?  Describe  the  most  common 
of  each.  If  you  had  sealed  up  the  solution  of  sugar,  would  crystals 
have  formed  ?  If  you  had  sealed  up  a  hot  solution  of  alum,  would  you 
expect  to  find  crystals  1  Try  these  experiments,  being  careful  that  the 
solutions  are  kept  air-tight. 

Experiment  13. — Latent  Heat  of  Solution. 

Materials  and  Apparatus. — Test-tubes,  basin,  thermometer,  ice,  salt. 

Mix  equal  parts  of  snow,  or  ice,  and  salt  in  a  basin.  Test  the  temper- 
ature from  time  to  time  with  a  thermometer  as  the  mass  is  stirred.  Does 
the  act  of  going  into  solution  consume  heat  ?  If  a  stirring-rod  is  made 
by  putting  a  small  test-tube  inside  a  larger  one,  with  the  space  between 
filled  with  water,  the  water  will  become  frozen.  The  heat  consumed  in 
changing  the  solids  to  solution  is  latent  or  insensible  heat.  If  crystalliza- 
tion is  the  opposite  of  solution,  should  we  .not  expect  that  latent  heat 
is  given  off?  Is  it  warmer  just  before  a  snow-storm  ? 


22  LABORATORY   MANUAL. 

Experiment  14.  —  Freezing  Mixture.     Latent  Heat,  Continued.    * 

(For  the  Teacher.)  —  Take  9  parts  by  weight  of  phosphate  of  sodium 
and  4  parts  by  weight  of  dilute  nitric  acid.  Cool  down  the  ingredients  by 
snow  and  salt  to  near  0°,  then  mix  the  ingredients  and  test  the  tempera- 
ture with  a  thermometer. 

How  low  a  temperature  have  you  reached  ?  How  i»  ice  made  arti- 
ficially 1  In  what  experiment  have  you  made  heat  latent  by  evapora- 
tion 1  Blow  upon  the  hand  with  the  mouth  wide  open.  Do  so  again 
Avith  the  lips  nearly  closed.  In  which  case  did  the  breath  seem  cold  1 
Why  was  it  cold  1  Does  condensation  produce  heat  ? 

Experiment  15.  —  Allotropism. 

Materials  and  Apparatus.  —  Test-tube,  beaker,  sulphur,  gas-lamp. 

Put  5  grams  of  sulphur  into  a  test-tube  and  melt  slowly  over  the 
gas-lamp  until  it  begins  to  boil.  Do  not  hold  the 
mouth  of  the  test-tube  too  near  the  flame  for  fear 
the  vapor  may  take  fire.  Notice  changes  of  color 
as  the  sulphur  melts.  Does  it  become  thin 
enough  to  flow  1  Does  it  when  about  to  boil  ? 
Pour  a  fine  stream  into  cold  water,  and  allow 
same  to  cool  slowly.  Are  these  two  parts  alike  1 
If  not,  how  do  they  differ  1  Examine  the  soft 
sulphur  after  two  or  three  days.  Has  it  become 
hard  and  yellow  ? 

When  a  substance  assumes  two  or  more  forms 

that  differ  only   in   physical  properties,   it   is   called   Allotropic.      The 
diamond,  charcoal,  coke,  and  lamp-black  are  allotropic  forms  of  carbon. 


.  —  Where  among  these  changes  lies  the  domain  of  chemistry?  All  the 
changes  we  have  made  suggest  two  distinct  groups  of  changes,  with  one  intermediate 
group. 

1st,  Those  changes  that  do  not  affect  the  nature  of  the  substance  ;  2d,  Those 
changes  which  do  affect  the  nature  of  the  substance  ;  and  3d,  Changes  that  seem 
to  have  affected  the  nature  of  the  substance,  but  in  reality  have  only  changed  the 
physical  state. 

Chemistry  is  concerned  directly  with  the  latter  two,  though  indirectly  with  the 
first-mentioned  class  of  changes. 

REVIEW  QUESTIONS. 

What  is  a  physical  change  ?  A  chemical  ?  A  physico-chemical  ?  Which 
of  the  experiments  you  have  tried  belong  to  physical  ?  Which  illustrate  chemi- 
cal and  which  physico-chemical  ?  Give  other  illustrations  of  each, 


24  LABORATORY    MANUAL. 

Experiment  16. — Solution  assists  Chemical  Change. 

Materials  and  Apparatus. — Mortar  and  pestle,  test-tubes,  bicarbonate 
of  soda,  tartaric  acid. 

Mix  in  a  dry  mortar  half  a  gram  each  of  dry  tartaric  acid  and 
bicarbonate  of  soda  (cream  of  tartar  and  baking-soda).  Does  a  chemical 
change  occur  ?  Now  make  solutions  in  test-tubes  of  each  of  the  same 
ingredients  and  pour  into  one  test-tube.  Does  a  chemical  change  occur  ? 
Hold  the  thumb  over  the  mouth  of  the  test-tube  for  a  moment.  Is  some 
gas  being  expelled  1  Why  is  there  chemical  change  in  one  case  and  not 
in  the  other  ?  Pour  water  into  the  mortar.  What  occurs  1  How  does 
solution  help  chemical  change  1 

Experiment  17. — Divisibility  of  Matter. 

Materials  and  Apparatus. — Purple  aniline,  a  litre  or  quart  flask. 

Into  a  litre  (1000  c.  c.)  of  water  in  a  flask  put  one  decigram  of  aniline 
powder.  Does  it  dissolve  ?  Pour  out  half  of  the  water  after  it  is 
thoroughly  mixed  and  add  as  much  clean  water.  How  much  aniline 
have  you  now  1  Pour  out  half  and  till  as  before.  How  much  aniline 
left  1  Repeat  this  as  long  as  you  can  see  a  trace  of  aniline  in  the  water, 
keeping  account  all  the  time  of  the  amount  of  aniline  remaining  in  the 
flask.  Divide  by  1000,  since  you  have  1000  c.  c.  of  water.  How  much 
aniline  have  you  in  each  c.  c.  1  Imagine  this  division  carried  on  until 
you  have  reached  the  smallest  possible  particle  of  aniline.  Call  it  a 
Molecule.  Thus  each  substance  may  be  conceived  to  be  made  up  of 
molecules.  In  how  many  ways  can  molecules  differ  1 

Experiment  18. — Porosity  of  Matter. 

Apparatus.  — Beakers,  air-pump. 

Fill  a  glass  with  cold  water  and  stand  it  in  a  warm  room  for  half  an 
hour.  Do  you  notice  air-bubbles  rising  ?  Why  does  air  come  from  the 
water  under  these  circumstances  ?  Put  a  fresh  glass  of  water  under  the 
receiver  of  the  air-pump  and  exhaust  the  air.  Can  you  pump  air  out  of 
water  1  How  can  a  vessel  full  of  water  contain  air  1  When  the  air  is 
exhausted  is  the  volume  of  water  less  1  When  alum  went  into  solution 
did  it  displace  part  of  the  water,  that  is,  increase  its  volume  ?  What 
must  be  the  shape  of  molecules  of  water  to  allow  small  particles  of  a  solid 
or  gas  to  fill  part  of  the  same  space  1  Could  you  illustrate  with  oranges 
and  peas  ?  AVill  a  c.  c.  of  water  saturated  with  alum,  hold  as  much 
sugar  in  solution  as  a  c.  c.  of  water  containing  no  alum  ?  Try  it. 


26  LABORATORY   MANUAL. 

Experiment  19. — Atoms. 

Materials  and  Apparatus. — Sugar,  platinum-foil,  gas-burner,  tongs. 

Weigh  out  a  decigram  of  sugar,  place  upon  the  platinum-foil,  and 
burn  it  over  the  flame. 

Is  the  charred  mass  lighter  than  a  decigram  1  What  is  it  1  What 
has  become  of  the  rest  of  the  substance  ? 

Is  this  a  chemical  or  a  physical  change  1 

If  you  had  been  able  to  take  a  molecule  of  sugar,  and  had  thus  burned 
away  part  of  it,  would  you  then  have  less  than  a  molecule  of  matter  ? 

Would  it  be  sugar  ?     How  do  you  know  *? 

What  force  divided  the  molecules  of  sugar  into  parts  of  different 
kinds  1 

Call  these  parts  of  a  molecule  Atoms. 

Experiment  20. — Compounds  and  Elements. 

(For  the  Teacher}  Materials  and  Apparatus. — Chemical  battery  U- 
tube,  sulphuric  acid. 

Decompose  water  and  develop  the  answers  to  the  following  questions 
with  others  that  may  arise : 

How  are  the  volumes  of  the  two  gases  that  arise  from  the  poles  of 
the  battery  related  1  Is  this  a  chemical  change  ? 

Is  the  same  kind  of  gas  found  in  each  tube  1 

Test  by  a  burning  splinter.     Do  they  behave  alike  ?       • 

What  must  be  the  proportion  of  each  of  these  gases  in  each  mole- 
cule of  water  ? 

If  a  substance  is  composed  of  two  or  more  simple  substances,  the  first 
is  compound  with  reference  to  the  constituents.  Hence  we  call  a  sub- 
stance that  the  chemist  is  unable  to  divide  into  two  or  more  substances 
an  element,  and  two  or  more  elements  combined,  a  compound. 

Where  would  you  class  each  of  the  following  : — sugar,  paper,  wood, 
sulphur,  iron  1 

NOTE. — The  table  on  page  28  gives  a  list  of  the  most  common  elements  that 
form  the  common  compounds,  though  some  new  process  may  be  discovered  by  which 
some  of  these  may  be  divided  into  two  or  more  other  elements.  New  elements 
may  be  discovered  from  time  to  time,  though  not  probably  of  great  practical  im- 
portance. 

The  symbols  are  used  for  convenience,  and  should  be  learned  as  fast  as  the 
elements  are  studied. 

Atomic  weights  are  simply  for  reference.  Symbols  marked  thus  (')  are  called 
monads;  ("),  diads ;  (;//),  triads;  (""),  tetrads.  (For  explanation  of  these  terms, 
see  Exp.  39. ) 


LABORATORY    MANUAL. 


SYMBOLS  AND  ATOMIC  WEIGHTS  OF  THE  COMMON  ELEMENTS. 


ELEMENT. 

SYMBOL. 

ATOMIC 
WEIGHT. 

ELEMENT. 

SYMBOL. 

ATOMIC 
WEIGHT. 

Aluminum 

Al"" 

27.04 

Lead  (Plumbum)    . 

Pb" 

206.4 

Antimony  (Stibium). 

SV"  ' 

119.6 

Magnesium   .... 

Mg" 

23.94 

Arsenic.     .... 

As"' 

74.9 

Manganese    .... 

Mn" 

54.8 

Barium       .... 
Bismuth     .... 

Ba" 

Bi"' 

136.9 
207.3 

Mercury  (Hydrargyrum) 
Molibdenurn.     .     . 

Hg" 
Mo" 

199.8 
95.9 

Boron    

B"' 

10.9 

Nickel  (?)      .... 

Ni" 

58.56 

Bromine     .... 

Br' 

79.76 

Nitrogen  

N'" 

14.01 

Cadmium  .... 

Cd" 

111.7 

Oxygen    

0' 

15.96 

Calcium      .... 

Ca" 

39.91 

Phosphorus  .... 

F" 

30.96 

Carbon  

C"" 

11.97 

Platinum  

Pt"" 

194.3 

Chlorine     .... 

Cl' 

35.37 

Potassium  (Kalium)     . 

K' 

39.03 

Chromium 

Cr"" 

52.  45 

Silicon 

Si"" 

28. 

Cobalt  (?)   .... 

Co" 

58.74 

Silver  (Argentum)  .     . 

Ag' 

107.66 

Copper  

Cn" 

63.18 

Sodium  (Natrium)  . 

Na' 

23. 

Fluorine     .... 

F 

19.06 

Strontium     .... 

Sr" 

87.3 

Gold  (Auruin) 

Au'" 

196.7 

Sulphur   

S" 

31.98 

Hydrogen  .... 

H' 

1. 

Tin  (Stannum)  . 

Sn"&"" 

117.4 

Iodine  

I' 

126.54 

Zinc     

Zn" 

65.1 

Iron  (Ferrum) 

Fe"&"' 

55.88 

Experiment  21. — Compounds,  Continued. 

How  can  so  few  elements  make  up  the  innumerable  compounds  that 
are  found  in  nature  ?  Take  the  four  digits,  1,  2,  3,  and  4.  Arrange 
them  in  all  the  different  ways  possible  to  form  numbers.  Thus,  1234, 
1324,  4321,  etc.  How  many  numbers  can  you  form1?  If  you  arrange 
them  in  all  the  possible  combinations  of  twos,  threes,  and  fours,  how 
many  numbers  can  you  form  ?  If  there  were  but  four  elements,  and  this 
were  the  only  way  they  could  unite,  how  many  compounds  would  there 
be  ?  How  many  words  in  the  English  language  formed  from  26  letters  1 
Has  the  limit  been  reached  ? 

Experiment  22. — Analysis. 

Materials  and  Apparatus. — A  hard  glass  test-tube,  red  oxide  of  mer- 
cury, a  long  splinter  of  pine,  gas-lamp. 

Put  a  gram  of  red  oxide  of  mercury  (HgO)  into  a  hard  glass  test-tube. 
Heat  it  slowly  over  the  gas-burner.  Thrust  a  glowing  splinter  down  the 
test-tube  from  time  to  time  to  determine  if  a  gas  like  one  of  those 
obtained  by  the  decomposition  of  water  is  being  given  off. 

Is  it  true  ?  Which  one  is  it  ?  Is  there  a  new  substance  found  in  the 
test-tube  1  What  does  it  look  like  ?  What  force  separated  oxide  of 
mercury  into  two  parts  1  Is  oxide  of  mercury  a  compound  ?  Of  what 
is  it  composed  ?  This  process  is  called  Analysis.  In  what  other  experi- 
ment have  you  used  this  process  ? 


30 


LABORATORY   MANUAL. 


Experiment  23. — Synthesis. 

Apparatus. — A  lamp  chimney  or  a  large  glass  tube,  gas-burner. 

Hold  a  cold,  dry  lamp-chimney  over  the  burning  gas-lamp  for  a  moment. 
Is  water  condensed  upon  the  inside  1  Where  did  it  come  from  1  Illu- 
minating gas  is  made  up  largely  of  hydrogen.  Air  contains  a  large 
quantity  of  oxygen.  When  these  two  unite  what  compound  do  they 
form  ?  In  wrhat  experiment  have  we  performed  the  opposite  of  this 
experiment  ?  Combining  elements  to  form  compounds  is  called  synthesis. 
In  what  experiment  have  you  used  this  process  before  ?  What  did  it 
illustrate  ?  What  force  did  you  employ  in  each  1 

Experiment  24. — Acids  and  Alkalies. 

Materials. — Litmus,  acids,  caustic  soda,  ammonium  hydrate. 

Put  a  drop  of  hydrochloric  acid  upon  a  piece  of  blue  litmus  paper. 
How  does  it  affect  it  1  Try  other  acids.  Do  they  all  have  the  same  or 
different  effects  1  Pat  a  drop  of  ammonium  hydrate  upon  the  spot  made 
by  the  acid.  Try  caustic  soda.  Do  caustic  soda  and  ammonium  hydrate 
behave  alike  ?  Try  acid  and  red  litmus,  then  try  ammonium  hydrate, 
sodium  hydrate,  calcium  hydrate,  etc.,  with  blue  litmus.  In  which  cases 
do  you  get  a  decided  result  ? 

What  law  can  you  state  from  these  experiments  as  to  the  difference 
between  acids  and  hydrates  ?  Hydrates  are  a  type  of  a  class  of  sub- 
stances called  bases.  How  can  you  tell  these  bases  from  acids  ? 

All  elements  may  be  classified  in  reference  to  their  power  of  forming 
bases  or  acids ;  some,  however,  may  form  both  a  base  and  an  acid.  (See 
list  of  acid-forming  and  base-forming  elements,  below.) 

Put  a  drop  of  sulphuric  acid  in  5  or  6  c.  c.  of  water.  Taste  it.  Do 
so  with  acetic  and  hydrochloric  acids.  What  taste  have  they  ? 

What  elements  have  all  acids  in  common  1  (See  Symbols.)  Define  an 
acid.  Call  all  the  compounds  you  find  that  turn  red  litmus  blue,  Alkalies. 


COMMON  BASE-FORMING  ELEMENTS. 

Al  Cr  Mn 

(NH<)*  Co  Mg 

Sb  Cu  Hg 

Ba  Au  Ni 

Bi  H  Ft 

Cd  Fe  K 

Ca  Pb  Ag 


Na 
Sr 
Sn 
Zn 


COMMON  ACID-FORMING  ELEMENTS. 


As 

B 

Br 

C 

Cl 

Cr 

Cyt 


*  This  compound  of  nitrogen  and  hydrogen  is  classed  with  elements  because  it  forms  a  base 
like  sodium  and  potassium.  The  radical  NH*  is  called  ammonium  in  distinction  from  the  gas 
NH3  called  ammonia. 

t  Cy  is  a  compound  of  carbon  and  nitrogen,  C2N2,  which  acts  like  an  element  in  forming 
cyanic  acid.  The  radical  is  called  cyanogen. 


32  LABORATORY    MANUAL. 

Experiment  25.— Neutralization,  Formation  of  Salts. 

Materials  and  Apparatus.  — Beaker,  evaporating  dish,  water-bath, 
sodium  hydrate,  hydrochloric  acid,  litmus. 

Slowly  drop  HC1  into  25  c.  c.  of  the  solution  of  sodium  hydrate,  as 
long  as  it  shows  chemical  action.  Now  test  with  litmus  and  add  a 
drop  of  acid  or  alkali  until  it  turns  the  color  neither  red  nor  blue. 
It  is  now  neutral,  so  far  as  you  can  perceive  by  the  litmus.  A  change 
must  have  occurred  in  the  chemical  nature  of  both.  If  so,  new  sub- 
stances may  be  looked  for.  Transfer  to  the  evaporating  dish  and  place 
upon  the  water-bath  over  the  gas-lamp  to  evaporate.  Taste  the  dry  sub- 
stance obtained.  What  is  it  1 

Complete  the  equation,  NaHO  -f  HC1  =  NaCl  -f  ? 

NaCl  contains  the  first  element  of  the  alkaline  hydrate  and  the  last 
element  of  the  acid.  What  is  the  symbol  for  common  salt  ?  Is  this  salt 
acid,  alkaline,  or  neutral  1  Try  the  following  and  see  if  you  can  neutralize 
and  obtain  a  salt :  H2S04  +  2NH4HO  =  ?  +  ?  Evaporate  as  before. 
Is  this  salt  like  the  first  ?  When  zinc  and  HC1  were  used  to  liberate 
hydrogen,  was  a  salt  formed  1  Define  a  salt. 

Is  zinc  capable  of  turning  red  litmus  blue  1  Is  it  capable  of  neutral- 
izing the  power  of  an  acid  ?  Drop  a  grain  of  zinc  into  a  test-tube  and 
cover  with  dilute  hydrochloric  acid.  Is  there  chemical  action  ?  Is 
hydrogen  given  off?  Allow  it  to  stand  ovar  night.  Is  there  any  acid 
property  left  1  Is  there  a  solid  substance  formed  differing  from  zinc  1 
Compare  with  chloride  of  zinc.  Call  all  elements  that  can  neutralize  an 
acid,  a  base. 

Are  all  alkalies  that  you  have  tried  capable  of  forming  the  base  of  a 
salt  1  It  will  be  noticed  that  when  the  base  is  a  monad  it  combines  equally 
with  HC1.  Thus,  HCl  +  Na  =  NaCl-^-H.  But  when  a  diad  base  com- 
bines with  HC1  we  must  take  two  molecules  of  the  acid.  Thus, 
Zn  +  2HCl  =  ZnCl2  +  2H,  since  zinc  is  a  diad  and  hydrogen  a  monad. 
Again,  if  sodium  Na  combines  with  H2S04  it  takes  two  atoms  of  Na  to 
form  the  neutral  salt,  since  there  are  two  atoms  of  H.  Thus,  H2S04  + 
2Na  =  Na2S04  +  2H. 

It  will  be  seen  that  the  number  of  hydrogen  atoms  standing  first  in 
the  symbol  shows  the  combining-  power  of  the  acid,  and  that  this  hydrogen 
is  simply  thrown  out  to  give  place  to  a  base  of  the  same  number  of  hands 
or  combining  powers.  Call  the  combining  power  of  an  acid  its  basicity. 
As  there  are  but  few  common  acids,  it  is  important  to  know  the  basicity 
of  each  so  you  may  use  them  readily  in  writing  symbols. 


34 


LABORATORY    MANUAL. 

TABLE. — BASICITY  OF  ACIDS. 


MONOBASIC  ACIDS. 

DIBASIC. 

TBIBASIC. 

TETRABASIC. 

Hydrochloric,  HCl. 

Sulphuric,  H.SO*. 

Phosphoric,  H3P(K. 

Hydroferrocyariic, 

Chloric,  HC103. 

Sulphurous,  H2S03. 

Arsenic,  H3AsO*. 

HtFeCy8. 

Hydroiodic,  HI. 

Hydrosulphuric,  H2S. 

Arsenious,  H3As03- 

lodic,  HI03- 

Carbonic,  H2C03. 

Hydroferrieyanic, 

Hydrobromic,  HBr. 

Chromic,  H2CrO«. 

H3FeCye. 

Nitric,  HN03. 

Oxalic,  H2C2(K. 

Nitrous,  HN02. 

Tartaric,  H2C.H*0.. 

Acetic,  HC2H302. 

The  above  table  shows  the  combining  power  or  basicity  of  each  of 
the  most  common  acids.  As  before  stated,  the  atom  or  atoms  of 
hydrogen  standing  first  in  the  symbol  may  be  thrown  out  by  a  base  to 
form  a  salt.  In  this  light  an  acid  may  be  regarded  as  a  salt  of  hydrogen, 
and  the  name  hydric  chloride  and  hydric  sulphate  used  instead  of 
hydrochloric  and  sulphuric  acid. 

The  group  of  atoms  following  hydrogen  in  the  symbol  of  sulphuric 
acid  (H2S04)  is  called  an  acid  radical.  S04  is  the  acid  radical  which 
combines  with  a  base  to  form  a  salt. 

All  salts  in  which  S04  is  found  will  be  called  a  sulphate.  Salts 
formed  from  H2C03  will  have  the  radical  C03,  and  are  called  carbonates. 
When  there  are  two  acids  formed  by  the  same  elements,  differing  only 
in  the  number  of  atoms  of  oxygen,  as  H2S04  and  H2S03,  the  name  of 
the  acid  containing  the  greater  number  of  oxygen  atoms  terminates  in 
ic  and  that  of  the  one  containing  the  less  in  OUS. 

When  the  name  of  the  acid  terminates  in  ic  the  salt  formed  ter- 
minates in  ate. 

When  the  name  of  the  acid  terminates  in  OUS  the  salt  formed 
terminates  in  ite. 

Thus,  nitrw  acid  forms  nitrates,  while  nitrous  acid  forms  nitrites.  If 
an  acid  contains  no  oxygen,  as  HCl,  its  salts  terminate  in  itle.  Thus, 
from  HCl,  chlorides;  from  HBr,  "bromides,  etc.  Regarding  an  acid  as  a 
salt  of  hydrogen,  H2S04  would  be  called  hydrogen  sulphate,  or  sulphate 
of  hydrogen,  or  hydric  sulphate.  Replace  the  hydrogen  by  a  base,  as 
sodium,  we  have 

H2S04  +  2Na  =  Na3S04  +  2H. 


The  salt,  according  to  the  above,  would  be  called  sodium  sulphate, 
sulphate  of  sodium,  or  sodic  sulphate. 


36 


LABORATORY    MANUAL. 


Experiment  26. — Indestructibility  of  Matter  or  Persistence  of  Mass. 

(For  the  Teacher.)— (a)  Bal- 
ance electrolysis  apparatus  on  scale- 
pan.  Decompose  the  water,  col- 
lecting the  gases.  As  this  chem- 
ical change  takes  place  is  there  a 
change  of  volume1?  Is  there  any 
change  of  mass  ?  Is  any  matter 
destroyed  ? 

(b)    Perform     experiment      in 
synthesis  of   water  in  eudiometer. 
Have   eudiometer   balanced  before 
and    after    passing   the   spark.      Is 
Fig.  6. — EUDIOMETER,  there  any  loss  of  mass  ? 

Experiment  27.* — Indestructibility  of  Matter,  Continued. 

(For  the  Pupil.} — Place  less  than  a  gram  of  mercuric  sulphocyanate  in 
an  open  dish  upon  the  balance-pan.  Counterpoise  it,  then  ignite  it  with 
a  hot  iron  or  a  match,  and  note  effect  upon  the  equilibrium  of  the  balance. 

Experiment  28. — Law  of  Definite  Proportions. 

(«)  Evaporate  a  few  drops  of  hydrochloric  acid  solution,  then  a  few 
drops  of  ammonia  solution. 

(b)  With  some  convenient  measure  (e.g.,  a  test-tube  with  a  slip  of 
paper  gummed  upon  its  side)  measure  out  two  equal  portions  of  strong- 
ammonia  water.     Neutralize  exactly  with  strong  hydrochloric  acid  solu- 
tion  (using   litmus    paper)    these    two    portions   added   together  in   an 
evaporating  dish,  the  weight  of  which  has  been  determined.     Note  care- 
fully the  amount  of  acid  used.     Evaporate  the  contents  of  the  dish  to 
dryness  on  the  water-bath  and  weigh  the  salt  obtained. 

(c)  Take  again  the  same  quantity  of  hydrochloric  acid  and  add  to  it 
only  one  measureful  of  ammonia  water ;  evaporate  as  before,  and  weigh 
the  salt  obtained. 

(d)  Take  again  the  same  quantity  of  acid  and  add  to  it  three  measure- 
fuls  of  ammonia  water;  evaporate  as  before  and  weigh  the  salt  obtained. 
What  is  the  ratio  between  the  quantities  of  salt  obtained  in  the  three 
cases  1     What  is  the  logic  of  this  experiment  ? 


*  Experiments  Nos.  27,  28,  and  29  were  prepared  by  Professor  A.  V.  E.  Young,  Professor  of 
Chemistry  in  the  Northwestern  University  at  Evanston,  111. 


38  LABORATORY   MANUAL. 

Experiment  29. — Law  of  Multiple  Proportions. 

NOTE.  — In  the  following  experiment  two  or  more  pupils  may  work  together. 

(a)  Take  iodine  and   mercury  in   the  proportions  of   254  parts  by 
weight  of  the  former  and  200  parts  of  the  latter.     For  the  actual  experi- 
ment weigh  out  accurately  (in  some  glass  vessel)   6.3  grams  of  iodine. 
Then  weigh  out  5  grams  of  mercury  and  transfer  it  to  a  mortar  previously 
weighed.     Add  two  -or  three  drops  of  alcohol  and  then  a  small  portion  of 
the  iodine.     Triturate  thoroughly  until  the  mass  is  dry.     This  should 
yield  a  red  powder.     A  drop  of  alcohol  added  now  should  be  but  slightly, 
if  at  all,  colored  by  the  iodine. 

If  the  powder  is  not  a  bright  red,  and  the  alcohol  is  colored  by 
iodine,  the  trituration  must  be  continued  with  perhaps  a  little  warming 
in  the  water-bath.  When  this  operation  is  finished,  weigh  mortar  and 
contents.  Then  take  a  small  portion  of  the  red  powder,  such  as  would 
be  taken  up  on  the  point  of  a  penknife,  transfer  it  to  a  test-tube,  and 
add  a  little  alcohol  and  boil.  Be  careful  that  the  vapor  of  alcohol  does 
not  take  fire  at  the  gas-jet.  The  powder,  or  a  portion  of  it,  is  dissolved 
by  the  boiling  alcohol  and  redeposited,  on  slow  cooling,  in  shining  crystal- 
line- scales  of  a  brilliant  scarlet. 

Sublime  another  small  portion  of  the  red  powder  in  a  dry  test-tube ; 
note  the  red  and  yellow  sublimate.  Rub  the  yellow  with  a  glass  rod ; 
note  the  effect.  The  substance  obtained  in  this  experiment  is  mercuric 
iodide,  symbol  HgI2,  soluble  in  hot  alcohol,  subliming  without  decompo- 
sition. The  red  and  yellow  sublimates  are  allotropic  forms  of  the  same 
composition.  By  gentle  friction  the  yellow  is  converted  into  the  red. 

(b)  Take  5.6  grams  of  the  mercuric  iodide  obtained  in  (a)  [equal  to 
half  the  yield  (6.3  plus  5)]  and  2.5  grams  of  mercury  [equal  to  half  the 
amount  used  in  (a)].    Add  the  mercury,  in  small  portions  at  a  time,  to  the 
mercuric  iodide  in  a  mortar  and  triturate  thoroughly,  using  considerable 
pressure.     This  should  yield  a  yellowish-green  powder.     If  the  color  is 
not  satisfactory,  continue  the  trituration  with  the  addition  of  a  few  drops 
of  alcohol.      If  still  unsatisfactory,  allow  it  to  stand  for  several  hours. 
When  finished,  weigh  mortar  and  contents.     Treat  small  portion  as  in  («) 
with  boiling  alcohol  [remember  caution].     Heating  another  portion  in  dry 
tube  turns  it  brick-red  and  yields  a  sublimate  as  in  («),  and  mercury. 

The  green  powder  thus  obtained  is  mercurous  iodide,  symbol  Hgl ; 
insoluble  in  hot  alcohol  and  decomposed  by  heat  into  mercury  and 
mercuric  iodide.  What  are  the  relative  proportions  of  mercury  and 
iodine  in  the  two  substances  ?  What  is  the  logic  of  this  experiment  ? 


40 


LABORATORY   MANUAL. 


Experiment   30. — Chemistry  of  Water. 

(For  the  Teacher.) — Trough  of  water,  metallic  sodium,  wire-gauze, 
bottle  for  collecting  gas,  piece  of  wire  or  iron 
tongs. 

Roll  up  a  piece  of  metallic  sodium 
the  size  of  a  pea  in  a  small  piece  of 
gauze.  Invert  a  bottle  full  of  water  in 
the  water-trough.  IVy  means  of  a  stiff 
wire  or  a  pair  of  tongs,  thrust  the  gauze 
containing  the  sodium  under  the  mouth 
of  the  bottle.  Sodium  will  rapidly  unite 

with  one  of  the  elements  of  the  water,  setting  free  the  other.  In 
Experiment  20  we  learned  that  water  is  composed  of  two  gases  ;  that 
one  burns  with  a  blue  flame,  hydrogen.  The  other  will  not  burn,  but 
makes  a  burning  splinter  burn  more  rapidly,  oxygen. 

Test  this  gas.  Which  is  it?  When  hydrogen  burns  what  compound 
is  formed  ?  Write  an  account  of  this  experiment  and  be  sure  to  answer 
all  of  the  above  questions. 

Experiment   31.— Distillation. 

(For  the  Teacher)  Materials  and  Apparatus. — Large  retort  or  distill, 
condenser,  platinum-foil,  artesian  well-water  if  possible,  or  any  mineral 
water,  gas-lamp. 

Distill  water  before 
the  class.  What  proof 
can  you  see  that  impuri- 
ties are  left  in  the  retort  ? 
Does  well-water  contain 
solid  matter  in  solution  ? 
Can  you  tell  by  the  clear- 
ness 1  Put  a  few  drops 
of  well  -  water  upon  a 
clean  platinum-foil.  Hold 
it  over  the  gas-lamp  so  it 
will  not  boil  but  evaporate 
rapidly.  Is  there  a  sedi- 
ment left  upon  the  foil  ? 
Try  the  same  with  distilled  water, 
making  solutions 


Fig.  8.-  DISTILLATION. 


Why  is  distilled  water  valuable  in 
Why  has  it  not  an  agreeable  taste  ?    Why  are  mineral 


waters  desirable  ?     How  do  they  acquire  these  properties  1 


42  LABORATORY   MANUAL. 

Experiment  32.— Water  of  Crystallization, 

Materials  and  Apparatus. — Glass  tube  four  inches  long,  crystals  of 
zinc  sulphate,  sugar,  splinters  of  wood. 

Seal  up  one  end  of  a  glass  tube  by  a  blow-pipe  and  gas-lamp.  Exam- 
ine some  crystals  of  zinc  sulphate  as  to  color,  luster,  etc.  Put  them  into 
the  test-tube  and  heat  the  lower  end  gently.  Do  the  crystals  change  ? 
If  they  contain  water  it  will  be  driven  out  and  condense  near  the  upper 
end  of  the  tube.  Empty  the  tube  when  satisfied  as  to  the  presence  of 
water;  dry  it,  and  when  cool,  introduce  sugar,  and  heat  as  before.  Does 
it  contain  water  1  Dry  the  tube,  and  when  cool,  try  wood,  a  fiber  of 
meat,  a  green  stem,  etc.  Water  in  crystals  is  called  water  of  crystalliza- 
tion. How  can  you  tell  the  amount  of  water  in  a  pound  of  granulated 
sugar  ?  A  pound  of  steak  ?  How  much  of  the  human  body  is  water  ? 

Experiment  33. —  Efflorescence  and  Deliquescence. 

Materials  and  Apparatus. — Two  watch-glasses,  sodium  sulphate,  and 
calcium  chloride. 

Put  some  crystals  of  sodium  sulphate  upon  a  watch-glass  and  leave 
exposed  to  the  air  for  an  hour  or  more.  Do  they  lose  water  as  the  zinc 
sulphate  did  by  heating?  Are  the  crystals  destroyed?  When  crystals 
lose  their  water  of  crystallization  by  exposure  to  the  air,  they  are  said  to 
effloresce.  In  the  same  way  expose  some  calcium  chloride  to  the  air  and 
leave  over  night.  How  has  it  behaved  ?  Where  did  the  water  come 
from  ?  This  is  said  to  deliquesce.  How  do  these  two  salts  stand  with 
respect  to  their  attraction  for  water  ? 

If  you  wished  to  dry  the  moisture  out  of  air,  what  could  you  take  to 
do  it  with  ?  What  precaution  is  needed  if  you  wish  to  keep  these  two 
salts  in  good  condition  for  future  use  ?  Why  ? 

Experiment  34. — Impurities  in  Drinking-water.     Organic. 

Materials  and  Apparatus. — Beakers,  sulphuric  acid,  solution  of  potas- 
sium permanganate,  samples  of  water  from  different  sources,  including 
one  from  a  ditch  or  pool,  and  distilled  water. 

Fill  the  beakers,  each  with  a  sample  of  water,  and  one  with  ditch- 
water  and  another  with  distilled  water.  Put  a  few  drops  of  sulphuric 
acid  into  each  and  then  enough  of  the  permanganate  solution  to  give  the 
water  in  each  a  deep  purple  tint,  as  nearly  alike  as  possible.  Set  in  a 
warm  place  for  an  hour.  Organic  matter  bleaches  out  the  color,  or  rather 
decolorizes  the  permanganate  solution.  Which  one  shows  most  and 
which  one  least  organic  matter  ?  Why  1  How  does  water  acquire 
organic  matter  ?  What  is  organic  matter  ? 


44  LABORATORY   MANUAL. 

Experiment   35. — Impurities  in  Drinking-water.     Chlorides   or 

Chlorine. 

Materials  and  Apparatus. — Beakers,  silver  nitrate,  nitric  acid,  am- 
monium nitrate,  samples  of  water,  distilled  water,  and  common  salt. 

Concentrate  by  boiling  50  c.  c.  of  drinking-water  to  25  c.  c.;  add  a 
few  drops  of  nitric  acid  and  then  a  few  drops  of  silver-nitrate  solution. 
Now  take  25  c.  c.  each  of  distilled-water  and  salt-water  solution,  in 
separate  beakers,  and  treat  in  the  same  way  with  acid  and  silver  nitrate. 
These  two  are  only  for  comparison,  as  distilled  water  will  contain  no 
chlorine,  and  salt  water  a  great  deal,  as  salt  is  composed  of  equal  volumes 
of  sodium  and  chlorine  (NaCl). 

A  chloride,  as  you  see  by  the  salt  solution,  becomes  milky  by  the 
addition  of  silver  nitrate.  Does  the  sample  of  drinking-water  from  the 
pond  1  Does  that  from  a  deep  well  1  Distilled  water  ?  Sewage  is  a 
source  of  chlorine  in  drinking-water,  and  such  water  should  be  avoided, 
unless  it  comes  from  the  presence  of  common  salt,  as  in  salt  springs,  in 
which  case  it  does  no  harm. 

Experiment  36. — Hydrogen. 

Materials  and  Apparatus. — Test-tube,  granulated  zinc,  hydrochloric 
acid. 

Put  a  few  grains  of  zinc  into  a  test-tube  and  cover  with  dilute 
hydrochloric  acid.  What  occurs  1  Is  it  a  chemical  change  ?  After  the 
action  has  continued  for  a  few  minutes,  apply  a  lighted  splinter  to  the 
mouth  of  the  test-tube.  What  occurs  1  Does  this  prove  that  chemical 
change  has  occurred  ?  Such  bubbling  is  called  effervescence.  What  is 
the  gas  ?  In  what  experiments  have  you  made  or  seen  it  made  before  1 
Try  zinc  and  dilute  sulphuric  acid.  Is  the  same  gas  given  off  ?  Again 
try  iron  and  sulphuric  acid.  Do  you  find  the  same  gas  ?  Since  zinc 
and  iron  are  both  elements,  could  the  hydrogen  have  come  from  them  1 
Look  at  the  symbol  on  each  of  the  acid  bottles,  HC1  and  H3S04.  What 
element  in  common  ?  To  express  this  in  chemical  language,  we  write  : 
2HCl  +  Zn=  ZnCl2+2H.  This  reads,  two  molecules  of  hydrochloric 
acid  and  one  atom  of  zinc  form  one  molecule  of  zinc  chloride  and  two 
atoms  of  hydrogen. 

Again,  H2S04  +  Zn  =  ZnS04  +  2H  reads,  one  molecule  of  sulphuric 
acid  and  one  atom  of  zinc  form  one  molecule  of  zinc  sulphate  and  two 
atoms  of  hydrogen.  Evaporate  the  latter  nearly  to  dryness  ;  set  aside 
for  twenty-four  hours,  and  see  if  you  have  crystals  of  a  substance  that  is 
neither  zinc  nor  sulphuric  acid.  Compare  them  with  zinc  sulphate. 


4G  LABORATORY   MANUAL. 

NOTE.  — A  number  before  a  molecule  multiplies  the  molecule  ;  hence,  multiplies 
each  of  the  elements  in  the  molecule.  Thus,  2HC1  means  two  atoms  of  hy- 
drogen and  two  of  chlorine.  A  number  placed  below  and  after  a  letter  multiplies 
that  element  only.  Thus,  ZnCl2  means  that  the  molecule  of  zinc  sulphate  con- 
tains one  atom  of  zinc  and  two  atoms  of  chlorine.  What  symbol  expresses  a  mole- 
cule of  water  ? 

Experiment  37. — Hydrogen,  Continued. 

Apparatus. — Same  as  in  the  last,  a  perforated  cork  to  fit  the  test- 
tube,  glass  tubing,  rubber  connections. 

Put  zinc  and  HC1  as  before  into  the 
test-tube,  after  arranging  your  appa- 
ratus as  in  the  cut,  except  that  you 
need  no  burner  under  the  test-tube. 
Fill  a  test-tube  or  a  bottle  with  water 
and  invert  in  the  water-pan  over  the 
end  of  the  delivery-tube.  The  gas  at 
first  will  be  impure.  Ignite  and  h'Jl 
again.  After  it  becomes  pure  it  burns 

quietly  at  the  mouth  of  the  bottle.     Fill  one  half  full  and  then  let  the 
water  out  and  the  air  in. 

Hold  the  mouth  of  the  bottle  downwards  and  ignite  the  gas.  Does 
it  explode  or  burn  quietly  1  Why  ? 

Fill  a  bottle  with  pure  hydrogen  and  thrust  a  burning  splinter  into 
the  bottle.  Does  it  burn  at  the  surface  or  at  the  end  of  the  splinter  1 
Does  hydrogen  support  combustion  ? 

What  do  we  mean  by  an  explosion  1  Has  hydrogen  an  odor  ? 
Color  1  It  is  not  poisonous. 

Experiment  38. — Hydrogen,  Continued. 

(for  the  Teacher.)  —  Generate  a  large  quantity  of  hydrogen,  and 
^  collect  in  gas-bag.  Test  the  lightness  by  making  soap-bubbles.  Why 
do  they  rise  1  Do  they  explode  when  -  ignited  ?  Are  they  pure  hydro- 
gen1? Pass  some  of  the  gas  through  a  U-tube  filled  with  chloride  of 
lime.  Have  the  gas  issue  from  a  fine  tube,  of  metal  if  possible,  and 
ignite, — "  Philosopher's  Lamp."  Why  do  we  pass  it  through  chloride  of 
lime  1  Hold  a  cold  bell-jar  over  the  burning  jet.  What  is  formed  upon 
the  cold  surface  ?  Will  hydrogen  burn  if  the  air  is  excluded  ?  Thrust 
a  small-mouthed  bottle  over  the  jet.  Save  a  gas-bag  of  hydrogen  for 
the  next.  How  are  balloons  made  to  rise  ?  Can  you  pour  hydrogen 
upward  from  one  bottle  to  another  ?  Give  the  result  of  these  experi- 
ments. 


48  LABORATORY    MANUAL. 

Experiment  39. — Synthesis  of  Water  by  Volume. 

(For  the  Teacher.) — Prepare  a  gas-bag  of  oxygen,  a  eudiometer  for 
exploding  gases,  a  pan  of  mercury,  and  a  battery  and  coil. 

Fill  the  eudiometer  with  mercury  and  invert  in  the  pan  of  mercury. 
Introduce  equal  parts  by  volume  of  each  of  the  gases  oxygen  and  hydro- 
gen from  the  gas-bags,  and  allow  them  to  mix.  Pass  a  spark  through  the 
wires.  What  is  the  volume  of  the  gas  remaining  1  Re  verse  the  tube 
and  test  the  gas,  and  see  which  gas  is  left. 

Try  two  volumes  of  0  and  one  of  H,  Then  two  of  H  and  one  of  0. 
Which  proves  to  be  the  best  proportion  1  Is  water  formed  by  the 
explosion  ?  How  does  this  agree  with  Experiment  20  1 

If  you  could  weigh  a  litre  of  hydrogen  accurately,  you  would  find  it 
about  .089578  grams,  and  a  litre  of  oxygen  1.429  grams.  How  many 
times  heavier  is  oxygen  than  hydrogen  ? 

According  to  the  law  of  Ampere,  equal  volumes  of  all  gases  contain 
the  same  number  of  molecules  under  like  conditions  of  temperature 
and  pressure.  That  being  true,  how  would  the  molecules  of  the  two 
gases  compare? 

It  may  also  be  shown  that  the  molecule  of  each  gas  contains  the  same 
number  of  atoms.  If  that  is  so,  will  the  atoms  compare  in  the  same 
ratio?  Taking*!!  as  the  standard,  as  it  is  the  lightest  element,  what 
would  be  the  atomic  weight  of  0  '?  (See  Table  of  Atomic  Weights  and 
Elements,  page  28.)  What  would  be  the  weight  of  a  molecule  of  water 
compared  with  an  atom  of  hydrogen  ?  Name  all  the  facts  shown  in  the 
symbol  H20. 

Since  one  atom  of  oxygen  has  the  power  of  combining  with  two  atoms 
of  hydrogen,  let  us  represent  oxygen  as  having  two  hands  with  which  to 
grasp  other  elements,  and  hydrogen  but  one  hand.  Then  one  atom  of 
oxygen  could  grasp  two  atoms  of  hydrogen,  and  the  hands  are  all 
occupied.  In  the  table  of  elements  and  atomic  weights  you  will  see  that 
these  hands  or  combining  powers  are  indicated  thus,  H',  0",  N"',  C"". 
Carbon  and  oxygen  in  a  satisfied  union  would  then  be  C02. 

Write  the  symbol  for  a  union  of  hydrogen  and  carbon  ;  nitrogen  and 
hydrogen.  Call  all  one-handed  elements  Monads ;  two-handed,  Diads ; 
three,  Triads;  four,  Tetrads.  To  aid  the  memory,  make  a  table  as  on 
p.  50,  putting  down  in  its  proper  place  each  -element  as  you  learn  its 
combining  power  or  atomicity,  or  valency,  as  it  is  called.  There  will  be 
only  about  36,  as  we  will  not  study  the  rare  elements.  Some  elements, 
as  iron,  must  be  placed  in  two  columns. 


50 


LABORATORY   MANUAL. 


MONADS. 

DlADS. 

TRIADS. 

TETRADS. 

H 

0 

N 

c 

Experiment  40. — Chemistry  of  Air.    Weight  of  Air. 

If  provided  with  the  apparatus,  weigh  a  litre  of  air. 
How  much  does  it  weigh  1 

What  apparatus  is  used  to  measure  downward  pressure  of  air  1 
much  has  this  been  found  to  be  upon  a  square  centimetre  ? 


How 


Experiment  41. — Oxidation  of  Lead. 

Place  a  small  piece  of  lead  in  a  porcelain  crucible,  and  find  the  weight 
of  the  crucible  and  its  contents  and  of  stirring-rod  placed  in  it. 

Apply  heat  to  crucible  until  lead  is  melted,  and  then  stir  with  the  rod 
until  contents  takes  on  the  form  of  a  powder. 

Weigh  the  three  articles  again,  as  the  rod  will  have  some  particles 
attached  to  it. 

Does  the  powder  weigh  more  or  less  than  the  original  piece  of  lead  ? 

Is  this  an  exception  to  the  law  of  persistence  of  mass  1 

Where  did  the  additional  matter  come  from  ? 

If  you  are  not  certain  in  regard  to  this  answer,  try  also  the  following 
experiment : 

Experiment  42. — Preventing  Oxidation. 

Weigh  small  piece  of  lead  in  crucible  containing  enough  borax  to 
protect  lead  from  the  air  when  melted.  The  pupil  should  not  stir  the 
contents  of  the  crucible  in  this  experiment.  Does  same  change  take 
place  in  lead  as  before  1  Why  not  ? 

Do  contents  increase  in  weight  during  this  experiment  as  before  ? 
Explain. 


52  LABORATORY   MANUAL. 

Experiment  43. — Nitrogen  of  Air. 

Place  a  piece  of  phosphorus,  the  size  of  a  pea,  on  an  iron  sand-bath 
and  float  on  water  in  pneumatic  trough.  Ignite  phosphorus  and  cover 

immediately  with  bell-jar  or  large 
wide-mouthed  bottle.  Note  level 
of  water  inside  same  as  that  outside 
at  beginning  of  experiment.  Bub- 
bles of  air  escape.  Explain. 

Describe  substance  forming  on 
inside. 

Allow  action  to  proceed  and  set 
apparatus  aside  until  contents  be- 
comes clear.  Note  level  of  water 
inside. 

Compare  action  of  oxygen  on  phosphorus  with  its  action  on  lead  in 
previous  experiments.  What  is  color  of  powder  in  this  case  ? 

What  has  become  of  this  powder  ?  If  enough  phosphorus  was  used, 
it  has  taken  up  all  the  oxygen.  If  so,  what  proportion  of  original  volume 
of  air  was  the  oxygen  1  How  do  you  know  this  1 

Experiment  44. — Nitrogen. 

Apparatus.— Make  a  U-tube  with  one  end  drawn  out  to  a  point, 
leaving  small  opening. 

Place  the  other  end  of  this  tube  up  into  the  jar  and  push  the  jar 
down  into  the  water,  forcing  gas  out  the  fine  opening.  Try  to  light  this 
gas.  Does  it  resemble  hydrogen  in  this  respect  ? 

Experiment  45.— Nitrogen,  Continued. 

Remove  tube,  and  cover  mouth  of  the  jar  under  water  with  a  piece  of 
glass  and  invert.  Lower  a  burning  taper  or  glowing  pine  splinter  into 
gas.  Does  it  promote  combustion  ?  Does  it  resemble  oxygen  in  this 
respect  ?  This  gas  is  called  Nitrogen. 

What  proportion  of  the  original  volume  of  air  was  nitrogen  1  (See  Ex- 
periment 44.) 

Experiment  46. — Per  Cent,  of  Nitrogen. 

(For  the  Teacher.) — Mix  100  volumes  of  air  and  50  volumes  hydrogen 
in  a  eudiometer  and  explode  as  in  synthesis  of  water.  Water  formed 
is  condensed  and  shrinkage  in  volume  noted  is  63 — i.e.,  volume  at 
beginning,  150;  after  exploding,  87  :  thus,  150  —  87  =  loss  of  63.  Of 
this,  of  course  -J  is  oxygen  and  f  hydrogen.  Hence,  A^-  =  21  parts  of 
oxygen  in  100  parts  of  air. 


54 


LABORATORY   MANUAL. 


Experiment  47. — Analysis  of  Air. 

By  Prof .  LcRoy  C.  Cooley,  Ph.D.,  in  "A   Guide  to  Elementary  Chemistry  for 
Beginners." 

"We  set  out  now  to  find  how  many  cubic  centimetres  of  nitrogen 
and  how  many  of  oxygen  and  carbon  dioxide  there  are  in  100  c.  c. 
of  air. 

To  do  this  we  will  imprison  a  vesselful  of  air,  and  then  run  into  it  a 
liquid  which  will  absorb  both  the  oxygen  and  the  carbon  dioxide  com- 
pletely, and  leave  the  nitrogen.  We  can  then  measure  the  nitrogen 
which  is  left,  and  we  can  find  out  how  much  there  was  of  the  other 
two  by  measuring  the  liquid  which  had  gone  into  the  tube  to  take 
their  place. 

Oar  Apparatus. — I  take  a  test-tube,  t  (Fig.  11),  to  hold  the  air.  A 
six-inch  tube,  |  inch  in  diameter,  will  do  ;  an  eight- inch 
tube  of  the  same  diameter  is  better.  The  rubber  stopper,  c, 
is  so  large  that  its  small  end  will  enter  the  tube  only  about 
a  half-inch.  It  has  two  holes  ;  to  close  one  I  have  a  solid 
rod  of  glass,  s ;  for  the  other,  a  glass  tube  reaching  just  a 
very  little  below  the  cork,  as  shown.  A  piece  of  thin 
rubber  tubing,  //.,  is  cut  about  six  inches  long.  There  is  a 
pinch-cock,  p,  by  which  its  walls  may  be  pinched  so  as  to 
close  it  completely.  F  is  a  small  glass  funnel. 

The  lower  end  of  h  I  stretch  over  the  tube  in  the  cork 
c,  and  its  upper  end  1  fix  over  the   stem  of  Ft  and  then  I 
place  the  funnel  in  the  clamp  of  the   support,  as  shown  in 
Fig.  12,  and  remove  the  rod  s. 

The  Liquid. — To  absorb  the  oxygen  and 
carbon  dioxide  gases  I  use  a  mixture  of  pyro- 
gallic  acid  and  potassium  hydrate. 

I  take  a  small  teaspoonful  of  the  solid  acid 
and  pour  on  it  10  c.  c.  of  water;  it  will  soon 
dissolve.  To  this  I  then  add  5  c.  c.  of  strong 
solution  of  potassium  hydrate,  and  at  once  pour 
it  into  the  funnel.  Next,  I  hold  the  dish  below 
the  cork  and  open  the  pinch-cock  p  a  moment,  to 
let  the  liquid  run  down  and  fill  the  tubes  com- 
pletely. I  carefully  take  off  the  drop,  which 
hangs  at  the  lower  end  of  the  tube  below  the 
cork,  with  a  piece  of  filter-paper.  Fig.  12, 


Fig.  11. 


1U. 


56 


LABORATORY    MANUAL. 


Fig.  13. 


I  press  the  tube  t  up  over  the  cork  until  the  joint  is  air-tight,  as 
seen  in  Fig.  13,  and  after  a  minute  I  put  the  rod  s  into  the  open  hole 
of  the  cork.      I  have  now  imprisoned  a  tubeful  of 
air;   none  can  get  out,  and  no  more  can  get  in. 

I  left  the  hole  in  the  cork  open,  because  if  it 
were  not  open  the  pressure  of  the  cork  would  crowd 
the  air  below,  and  there  would  be  too  much  in  the 
tube ;  and  then,  too,  handling  the  tube  warmed  it, 
and  the  volume  of  air  changes  with  heat.  With 
the  hole  open,  the  air  in  the  tube  soon  conies  to 
be  just  as  warm  and  just  as  much  pressed  as  the 
air  outside.  Whenever  a  gas  of  any  kind  is  to  be 
measured,  its  temperature  and  pressure  must  be  the 
same  as  those  of  the  air  outside, 

The  Absorption. — I  now  press  the  pinch-cock^; 
a  little  stream  of  the  liquid  falls  into  t  at  once,  and 
then   drops   follow,  or,  if  the  tube  be   slightly  in- 
clined, a  slender  stream  will  flow  down  its  side.      It  will  continue  to 
enter  as  long  as  there  is  any  oxygen  or  carbon  dioxide  for  it  to  absorb, 
and  then  stop.     The  gas  which  is  left  in  the  tube  is  nitrogen. 

But  this  gas  is  crowded  down  by  the  pressure  of  the  liquid  in  the 
rubber  tube  and  funnel  above,  and  so  I  take  hold 
of  the  cork  c,  and  the  rim  of  t,  not  to  warm  the    j   j 
gas  with  my  hand,  and  lift  the  tube  bottom  up,  as    I 
shown  at  T  in  Fig.  14,  making  the  level  of  the    j 
liquid  the  same  in  the  tube  and  in  the  funnel.     I    { 
then  open  the  pinch-cock.      Some  of  the  liquid  will 
run  out  of    T.     When  the  liquid  stands  at  the 
same  level  in  the  tube  and  in  the  funnel,  I  close 
the  cock  and  bring  the  tube  down  again. 

The  almost  black  liquid  in  t  has  now  taken 
out  all  the  oxygen  and  carbon  dioxide  from  the 
tubeful  of  air,  and  left  all  the  nitrogen. 

The  Measuring. — I  must  measure  the  liquid  in 
the  tube  to  find  how  much  oxygen  was  taken  out — 
and  carbon  dioxide  also ;  but  the  volume  of  the 
carbon  dioxide,  in  so  small  a  quantity  of  air  as  we 
use,  is  so  little  that  we  cannot  measure  it  with  our  apparatus,  and  there- 
fore leave  it  out  of  account  in  this  experiment — -and  the  space  above  it 
to  find  how  much  nitrogen  was  left.  To  do  this  I  slip  two  small  rubber 
rings  upon  the  tube,  and  make  the  upper  edge  of  one  mark  the  place  of 


Fig.  14. 


58  LABORATORY    MANUAL. 

the  lower  end   of   the  cork,  and  of  the   other,  the   top   of  the  liquid. 
These  rings  must  not  afterward  Le  disturbed. 

I  may  now  remove  the  cork,  empty  the  tube,  rinse  it  with  water,  and 
then  let  the  last  drop  of  water  drain  away.  Finally,  I  use  my  graduated 
cylinder  to  find  out  exactly — 

How  many  c.  c.  of  water  will  fill  the  tube  to  the  first  ring  ? 
How  many  c.  c.  from  the  first  to  the  second  ring  ? 
The  Calculations. — From  these  two  numbers  we  can  find  what  part 
of  the  air  is  nitrogen  and  what  part  is  oxygen,  for  they  help  us  to  answer 
the  following  questions,  in  their  order,  as  shown  by  an  example  below : 
How  many  c.  c.  of  air  were  in  the  tube  at  first  ? 
How  many  c.  c.  of  nitrogen  did  this  air  yield  ? 
How  many  c.  c.  of  oxygen  did  the  same  air  yield  ? 
Then  what  fractional  part  of  the  air  is  nitrogen  ? 
What  fractional  part  of  the  air  is  oxygen  ? 
And  how  many  c.  c.  nitrogen  in  100  c.  c.  of  air  ? 
How  many  c.  c.  of  oxygen  in  100  c.  c.  of  air  1 
An  Example. — In  an  actual  experiment  it  was  found  to  take  of 
Water  to  fill  the  tube  to  the  first  ring,         .         .         .         .         6. 0  c.  c. 
Water  to  fill  the  tube  from  the  first  to  second  ring,      .         .       23.5  c.  c. 
Hence  the  number  of  c.  c.  of  air  taken,        ....       29.5  c.  c. 
And  the  number  of  c.  c.  of  nitrogen  found,  .         .         .       23.5  c.  c. 

And  the  number  of  c.  c.  of  oxygen  found,    .         .         .         .         6.0  c.  c. 

Now  this  would  show  plainly  that  |-||  of  the  air  is  nitrogen  and  j6^g 
of  it  is  oxygen.  Then  in  100  c.  c.  of  air  there  would  be 

Nitrogen, 79.66c.  c. 

Oxygen, 20.34  c.  c. 

Experiment  48. — Moisture  in  Air. 

Does  air  contain  any  other  gas  besides  N  and  0  ?  Breathe  into  a  cold, 
clean  flask.  What  is  found  on  the  inside  1  Where  does  this  come  from  ? 
Is  there  some  of  this  substance  in  the  air  all  of  the  time  ?  To  answer 
this  question,  think  of  the  collecting  of  water  on  sides  of  ice-pitcher  in 
summer.  Explain  this.  What  is  dew  ?  What  is  frost  ? 

Experiment  49.— Carbon  Dioxide  in  Air. 

Put  some  lime-water  into  the  flask,  breathe  into  it  again,  shake  up 
contents  of  flask  and  note  result,  or  use  beaker  of  lime-water  and  blow 
into  it  through  glass  tube.  The  gas  which  produces  this  result  is  called 
Carbon  Dioxide.  Where  does  it  come  from  in  this  experiment  ?  How 
many  substances  have  you  now  seen  to  be  in  the  air  ?  Is  their  proportion 
always  the  same  1  Is  air  a  chemical  compound  or  a  mechanical  mixture  ? 


60 


LABORATORY    MANUAL. 

Experiment  50. — Preparation  of  Oxygen. 


Fig.  15.— PREPARATION  OF  OXYGEN. 

Make  a  quantity  of  oxygen  as  directed  in  Eliot  and  Storer's  "Element- 
ary Manual  of  Chemistry,"  Experiment  4,  page  9. 

Experiment  51. — Characteristics  of  0. 

(For  the  Teacher.} — Let  the  teacher  perform  the  experiments  given 
in  the  text-book,  or  others,  if  he  prefers,  to  illustrate  the  properties  of 
oxygen.  Pupils  should  note  each  step  in  each  experiment  and  write  out 
carefully  what  is  shown  by  it. 

Experiment  52.— Effect  of  0  on  Blood. 

N.B. — The  teacher  should  obtain,  from  the  nearest  slaughter-house,  a  quantity 
of  freshly-drawn  blood,  which  has  been  whipped  to  take  out  the  fibrine  so  that  it 
will  not  coagulate.  This  may  be  kept  for  several  days,  if  necessary,  in  a  cool  place, 
in  a  well-corked  bottle. 

Take  small  quantity  of  the  blood  in  a  test-tube.  Notice  its  dark-red 
color.  This  shows  you  the  color  of  so-called  venous  blood. 

Now  shake  up  the  test-tube,  placing  the  thumb  over  the  top. 

After  mixing  air  in  tube  with  the  blood,  remove  the  thumb  and  admit 
more  fresh  air,  and  shake  as  before. 

Repeat  this  several  times,  and  note  change  in  color  of  the  blood. 

Where  in  our  bodies  does  this  change  of  color  take  place  ?  What 
part  of  the  blood  gives  it  its  red  color  2 

Explain  the  change  which  takes  place  in  the  blood  in  the  lungs. 

What  change  takes  place  in  the  air  which  goes  to  the  lungs  ? 


62  LABORATORY   MANUAL. 

Experiment  53. — Oxygen  from  Green  Plants. 

(For  the  Teacher.) — A  tall,  wide-mouthed  bottle,  water-pan,  water- 
plants. 

Fill  a  tall,  wide-mouthed  bottle  half  full  of  water-plants.  The  light- 
green  fibrous  alga  spirogyra  is  the  best.  Fill  the  bottle  with  water  and 
invert  in  the  water-pan,  and  support  it  so  there  can  be  circulation  of  water 
from  the  bottle  to  the  pan.  Place  the  apparatus  in  the  sun-light  and 
allow  it  to  stand  until  a  sufficient  quantity  of  gas  is  generated.  Cover  the 
mouth  of  the  bottle  and  invert.  Test  for  oxygen  with  a  burning  splinter. 
Where  did  the  bubbles  of  gas  come  from  'J  What  had  the  light  to  do 
with  it  ?  What  effect  have  water-plants  upon  stagnant  pools  1  Do  plants 
in  the  air  do  the  same  work  1  What  is  the  gas  in  the  water  that  chloro- 
phyll decomposes  1  How  does  it  come  there  1  How  are  plants  and 
animals  related  in  respect  to  the  use  of  oxygen  1  Could  the  number  of 
plants  and  animals  in  an  aquarium  be  so  balanced  that  each  would  furnish 
the  gaseous  food  required  by  the  other  1 

Experiment  54. — Nitrous  Oxide. 

Compounds  of  N  and  0. — Place  a  small  quantity  of  ammonium 
nitrate  (NH4N03)  in  a  test-tube  and  heat. 

Hold  piece  of  cold  glass  near  mouth  of  tube,  and  note  what  collects 
upon  it. 

What  is  the  first  change  which  takes  place  in  the  contents  of  the 
tube  1  What  is  the  next  ?  Which  is  chemical  ?  Is  there  any  residue  1 

Complete  the  equation,  NH4N03  =  2H20-K 

Experiment  55.— Nitrous  Oxide,  Continued. 


Fig.  16.— PREPARING  NITROUS  OXIDE. 


Arrange  apparatus  as  shown  in  Fig.  16. 

Test  gas  collected  in  bottle  with  glowing  splinter,  as  in  experiment 
with  oxygen.     What  is  the  result  1 


G4 


LABORATORY    MANUAL. 


Experiment  56.  — Nitric  Oxide. 

N.B. — To  be  performed  under  the  hood  or  where  there  is  a  good  draught. 

Iii  flask  used  for  generating  hydrogen  place  a  few  grams  of  sheet- 
copper  cut  in  small  pieces,  and  connect  delivery-tube  to  wide-mouthed 
bottle  in  pneumatic  trough  as  in  preparing  hydrogen.  Pour  in  enough 
50  per  cent,  nitric  acid  to  cover  the  copper  and  to  come  above  the  opening 
of  the  thistle-tube.  (Do  not  inhale  the  fumes.) 

Collect  the  gas  given  off,  by  displacement  of  water,  as  in  the  case  of 
hydrogen.  What  is  in  the  generating  flask  before  the  beginning  of  the 

experiment  ?  What  is  the 
color  of  the  gas  given  off  at 
first  1  What  is  the  color 
after  a  quantity  has  been 
collected  over  water  ? 

NOTE.  —  As  soon  as  a 
small  quantity  of  the  gas  has 
been  collected,  prevent  further 
formation  of  it,  by  disconnect- 
ing the  generating  flask,  and 
pouring  off  liquid  which  it 
contains,  into  an  evaporat- 
ing dish  and  setting  it  aside. 
Rinse  copper  quickly  in 
plenty  of  water  and  put  away 
Fig.  17.— PREPARING  NO,  to  be  used  again. 

Experiment  57.— Nitric  Oxide,  Continued. 

Allow  a  small  amount  of  the  gas  which  you  have  collected  to  escape 
into  the  air.  What  is  the  color  when  it  comes  in  contact  with  the 
air?  What  element  in  air  is  most  active  and  most  likely  to  be  the 
cause  of  this  change?  Does  the  gas  after  it  has  passed  into  the  air 
resemble  that  in  the  generating  flask  when  action  first  commenced  ? 
How  do  you  account  for  this  ?  Complete  the  equation, 

3Cu  +  8HN03  =  3Cu(N03)  +  H20  +  NO. 

This  represents  the  chemical  action  in  making  nitric  oxide  (NO). 
NO-f-0  =  N02   (nitrogen  peroxide). 

How  many  oxides  of  nitrogen  are  described  in  your  text-book  ? 
Make  a  table  of  them,  showing  relation  of  N  to  0,  by  weight  in  each. 

What  fundamental  law  of  chemical  action  may  be  derived  from  a 
consideration  of  these  compounds  ? 

What  experiments  have  we  made  use  of  in  reaching  the  same  law  ? 

Evaporate  a  small  portion  of  the  liquid.  Set  aside  in  evaporating 
dish.  What  remains  ?  Name  of  the  substance  ?  Its  formula  1 


66 


LABORATORY    MANUAL. 


Fig.  ^.-PREPARING  NH3. 


Experiment  58.—  Compound  of  N  and  H. 

Mix  five  grams  of  ammonium  chloride  and  same  weight  of  cold, 
freshly  slaked  lime,  Ca(OH)2.  Arrange  flask 
as  shown  in  cut.  Note  odor  of  gas  given  off. 
Warm  gently  for  half  a  minute  and  try  to  col- 
lect gas  by  displacement  of  water  as  in  case  of 
NO  and  H.  Do  you  succeed  in  thus  collecting 
the  gas  ?  Test  the  water  with  litmus.  Is  the 
liquid  acid  or  alkaline  1  What  is  the  liquid  ? 

Experiment  59.—  Ammonia  from  Ammonium 
Chloride. 

Note  odor  of  ammonium  chloride.    Mix  small 
quantity  with  lime  as  in  preceding  experiment. 
What  is  the  odor1?     This  substance  is  called 
ammonia  (NH3).     The  complete  equation  is, 
2NH4C1  H-  Ca(OH)s  =  2NH3  +  CaCl2  +  2H20, 

also  NH3  +H20  =  NH4HO. 
Put  some  solution  of  ammonium  chloride   in  a  test-tube  and  add  a 
few  drops  of  KOH.     Result  ?     Write  the  equation. 

Experiment  60.—  Chlorine. 

Materials  and  Apparatus.  —  Test-tube,  pulverized  metallic  antimony, 
black  oxide  of   manganese,  hydrochloric  acid,  piece  of  calico. 

Put  a  gram  of  oxide  of  manganese  into  a  test-tube  and  cover  with 
HC1.  Heat  gently,  being  careful  not  to  breathe  any 
gas  that  may  be  given  off.  What  is  the  color  of 
the  gas  ?  Hold  a  piece  of  wet  calico  over  the  mouth 
of  the  test-tube  so  the  escaping  chlorine  will  have 
to  pass  through  it.  How  does  it  affect  the  color  1 
Try  dry  calico.  Write  upon  a  piece  of  paper  with 
carmine  ink  and  hold  it  over  the  test-tube.  Is  it 
bleached  ?  Try  common  print  moistened.  Why 
is  it  not  bleached  ?  Incline  the  tube  and  see  if 
the  gas  falls  or  rises  as  it  escapes.  Drop  some 
powdered  antimony  into  the  tube  as  the  gas  is 
rising  freely.  What  occurs  1 

Complete  the  equation,  Mn02  +  4HC1  =  MnCl2  +  2H20  +  ? 

Again,  Sb-f3Cl  =  ?      What  is   the   cause  of  the  fire  seen  in  the 
test-tube  1     Of  what  is  the  white  smoke  composed  ? 


68  LABORATORY    MANUAL. 

Experiment  61.—  Chlorine,  Continued. 

(For  the   Teacher.)  —  Let  the   teacher   prepare   chlorine  in  a 
quantity,  and  apply  the  following  tests.     Fill  several  bottles  with  the  gas 
and  pass  the  gas  through  water  in  Wolfe  bottles  to  prepare  chlorine  water. 

Wet  a  strip  of  paper  in  warm  water  and  thrust  it  down  into  a  bottle 
of  gas.  Does  it  show  chemical  action  ? 

Lower  a  lighted  taper  into  a  bottle  of  gas.     Will  it  burn  1 

Light  a  ribbon  of  manganese  and  lower  it  into  a  bottle  of  gas. 
Will  it  burn  ? 

Set  a  bottle  of  chlorine  water  inverted  in  a  pan  of  chlorine  water 
in  the  sun-light.  Examine  after  twenty-four  hours.  Has  gas  col- 
lected ? 

What  is  it  ?     Where  did  it  come  from  1 

Insert  a  cork,  with  a  short  piece  of  tubing  through  it,  into  a  bottle 
full  of  chlorine.  Invert  this  over  a  beaker  of  water.  After  half  an  hour 
or  so  examine.  Why  has  water  partly  filled  the  bottle  ? 

Put  10  grams  of  chloride  of  lime  into  a  litre  of  water;  add  a  few 
drops  of  HC1. 

Immerse  calico,  etc.,  to  see  if  it  bleaches  readily.  Will  chlorine 
water  remove  fruit  stains  1 

Experiment  62.—  Chlorine  Water. 

Drop  a  few  crystals  of  potassium  chlorate  (KC103)  into  a  test-tube  and 
add  hydrochloric  acid.  Warm  gently,  and  when  the  yellowish  gas  begins 
to  appear,  add  4  or  5  c.  c.  of  water. 

As  fast  as  chlorine  is  set  free  it  will  be  absorbed  by  the  water.  The 
reaction  is 

4HC1  +  2KC103  =  2KC14-2H20  +  C1204  +  2C1. 

What  effect  has  this  water  upon  vegetable  coloring-matter,  as  litmus, 
carmine,  etc.  1 

In  nearly  all  cases  where  bleaching  is  done  by  chlorine  the  reac- 
tion is 


The  oxygen  set  free  burns  up  the  coloring-matter,  so  that  in  reality 

oxygen  does  the  bleaching. 

In  what  form  is  chlorine  sold  by  the  druggist  for  bleaching  1 

Write  a  full  account  of  the  preparation  of  chlorine,  giving  its  chief 

characteristics,  and  the  uses  you  see  that  can  be  made  of  it, 


70  LABORATORY   MANUAL. 

Experiment  63. — Bromine. 

Materials  and  Apparatus. — Test-tube,  bromide  of  potassium,  man- 
ganese dioxide,  sulphuric  acid,  chlorine  water. 

Mix  half  a  gram  of  potassium  bromide  and  a  gram  of  manganese 
dioxide  in  a  test-tube,  and  cover  with  dilute  sulphuric  acid.  Heat  it.  Do 
not  breathe  the  gas.  How  does  bromine  differ  from  chlorine  1  Will  it 
bleach  calico  or  litmus  paper?  Compare  the  strength  of  chlorine  and 
bromine  by  the  following :  Make  a  solution  of  a  few  crystals  of  KBr  in 
two  or  three  c.  c.  of  water ;  add  a  drop  or  two  of  chlorine  water.  Is  Br 
set  free  ?  Explain  the  reaction,  KBr  +  Cl  =  KC1  +  (free)  Br. 

Has  chlorine  more  intense  chemical  action  than  bromine  ? 

Experiment  64. — Iodine. 

Materials  and  Apparatus. — Test-tube,  crystals  of  iodine,  potassium 
iodide,  alcohol,  chlorine  water,  starch  or  flour. 

Examine  a  crystal  of  iodine.  Does  it  stain  the  lingers  1  Dissolve  in 
water.  Make  a  thin  paste  of  starch  or  flour  and  add  a  drop  of  the  solu- 
tion. What  occurs  1  Dissolve  half  a  gram  of  potassium  iodide  (KI)  in 
two  or  three  c.  c.  of  water.  Add  to  this  a  c.  c.  of  starch  paste.  Does  it 
stain  as  the  element  iodine  ?  Why  not  1  Now  add  some  chlorine  water. 
Does  it  now1?  Complete  the  equation,  KI  +  C1=:KC1  +  ?  Does  this 
show  that  chlorine  is  stronger  than  iodine  1  Does  it  bleach  ?  How 
can  you  detect  the  presence  of  starch?  How  can  you  distinguish  be- 
tween these  three  elements  ?  How  are  they  related  ?  Dissolve  a  crys- 
tal of  iodine  in  alcohol.  This  is  a  tincture.  Allow  it  to  evaporate. 
Does  iodine  recrystallize  ? 

Experiment  65.— Hydrochloric  Acid,  HC1. 

Materials  and  Apparatus. — Test-tube,  litmus  paper,  common  salt, 
silver  nitrate,  ammonium  hydrate,  perforated  stopper  and  tubing,  beaker. 

Put  2  c.  c.  of  common  salt  into  a  test-tube  and  add  2  c.  c.  of  concen- 
trated sulphuric  acid.  Is  a  gas  given  off  ]  Test  with  a  strip  of  moistened 
blue  litmus.  Is  it  an  acid?  Complete  the  reaction,  2NaCl  -f  H2S04  = 
Na2S04  +  ?  Hold  an  open  ammonium-hydrate  bottle  near  the  mouth 
of  the  test-tube.  What  occurs  1  Now  hold  the  ammonium-hydrate  and  hy- 
drochloric-acid bottles  unstoppered  near  each  other.  Does  the  same  occur  ? 

HC1  +  NH3  =  NH4C1.  What  is  the  difference  between  the  reagent 
HC1  and  the  gas  you  have  generated  ?  Insert  the  stopper  and  pass  the 
gas  into  a  beaker  or  bottle  of  water  for  some  time.  Does  this  water 
acquire  acid  properties?  Is  sodium  sulphate  left  in  the  test-tube  after 
the  bubbling  ceases  ?  What  series  of  salts  is  formed  from  HC1  ? 


72  LABORATORY    MANUAL. 

Experiment  66.— Test  for  HC1. 

Take  two  or  three  drops  of  AgN03  in  a  test-tube,  and  add  a  few 
drops  of  HC1,  and  we  have  a  white  precipitate  of  silver  chloride,  thus, 
AgN03  +HC1  =  AgCl  +  HN03. 

Now  add  ammonium  hydrate  and  it  will  all  dissolve. 

Now  take  any  salt  of  hydrochloric  acid,  as  NH4C1,  CaCL,  or 
BaCl2,  and  add  AgN03. 

Do  you  get  the  same  white  precipitate  of  silver  chloride  1 

Try  its  solubility  ir>  NH4HO.  By  this  means  you  can  tell  any 
chloride. 

Experiment  67.— Nitric  Acid,  HN03. 

Let  the  teacher  prepare  nitric  acid  as  in  Eliot  and  Storer's  "Ele- 
mentary Chemistry,"  page  39,  Experiment  22. 

Pupils  try  the  following  characteristic  test  for  nitric  acid  and 
nitrates  : 

Make  a  solution  of  a  c.  c  of  ferrous  sulphate  (FeS04)  in  10  c.  c.  of 
water.  Add  to  this  2  c.  c.  of  H2S04  and  allow  to  cool. 

What  caused  the  heat  1 

Now  incline  the  test-tube  so  that  you  can  pour  in  two  or  three 
c.  c.  of  dilute  nitric  acid  without  its  mixing  with  the  heavy  solution 

below. 

« 

Where  the  two  liquids  meet,  a  brown  ring  will  be  formed. 

Now  do  the  same  again,  except  that  instead  of  using  nitric  acid  use 
a  solution  of  any  salt  of  nitric  acid,  as  KN03  or  NaN03. 

Does  it  act  in  the  same  way  ? 

By  this  test  you  may  tell  the  presence  of  nitric  acid  or  a  nitrate. 
Describe  the  manufacture  of  HN03. 

Experiment  68.— Sulphuric  Acid,  H2S04. 

The  manufacture  of  H2S04  is  too  difficult  to  perform  well.  Charac- 
teristic test  : 

Add  Bad  2  to  sulphuric  acid,  and  a  heavy  white  precipitate  of  BaSO , 
will  be  formed. 

Add  aqua  regia — a  mixture  of  concentrated  HC1  and  HN03 — and 
it  will  not  dissolve. 

Try  any  sulphate,  as  FeS04  or  CuS04,  and  see  if  they  act  the  same 
as  H2S04. 

Complete  these  equations  : 

BaCl2  +  H2S04  =  BaS04  -f  ?    BaCl3  +CuS04  =  ?  -h  ? 


74  LABORATORY   MANUAL. 

Experiment  69,  -Hydrosulphuric  Acid,  H2S. 

Place  a  small  fragment  of  ferrous  sulphide  (FeS)  in  a  test-tube  and  add 
4  or  5  c.  c.  of  water.  Now  add  a  c.  c.  of  sulphuric  acid. 

As  chemical  action  begins  notice  the  odor  of  rotten  eggs.  This  acid 
is  a  gas,  but  may  be  passed  into  water  and  the  saturated  water  used  as 
with  HC1. 

Complete  the  equation,  FeS  +  H2S04  =  H2S  +  ? 

Moisten  a  piece  of  filter-paper  with  lead  acetate  and  hold  it  over  the 
mouth  of  the  tube. 

Is  it  blackened  ? 

Do  not  allow  it  to  generate  gas  very  long,  as  it  is  poisonous.  Rinse 
out  the  tube,  saving  the  FeS  for  future  use. 

Characteristic  Test. — -Now  take  a  solution  of  a  salt  of  this  acid,  as 
ammonium  sulphide,  (NH4)2S,  in  a  test-tube.  Add  HC1. 

Is  H2S  given  off? 

Test  with  lead  acetate  as  before. 

Not  all  sulphides  will  be  shown  by  this  test ;  but  if  first  heated  on 
charcoal  with  Na2C03  and  then  heated  with  HC1,  they  yield  H2S. 

Can  you  generate  H2S  with  HC1  and  FeS  1     Try  it. 

For  generating  H2S  on  a  larger  scale,  see  Eliot  and  Storer's  "Ele- 
mentary Chemistry,"  page  76. 

Experiment  70.— Carbonic  Acid,  H2C03. 

This  acid  does  not  exist  uncombined,  as  it  splits  up  into  two  parts, 
thus  : 

H2C03  =  H20  +  C0o. 

Its  salts  are  abundant,  and  are  detected  by  the  following  :  * 

Take  any  carbonate  in  a  test-tube  and  add  HC1,  and  rapid  efferves- 
cence will  occur. 

While  the  test-tube  is  filling  with  gas,  fill  another  test-tube  J  full 
of  lime-water,  CaH202. 

Incline  the  test-tubes  as  if  you  would  pour  a  heavy  gas  from  the 
first  tube  into  the  lime-water.  After  a  moment  cover  the  mouth  of  the 
latter  tube  with  the  thumb  and  shake  thoroughly. 

Does  the  water  become  clouded  ?     What  gas  must  it  be  ? 

Characteristic  test  of  carbonates,  except  that  oxalates  may  yield  C02. 
The  oxalates,  however,  when  heated  on  a  platinum-foil  leave  a  black 
mass  and  carbonates  do  not. 

NOTE. — Let  the  teacher  give  several  salts  of  the  acids  given  here  to  be  distin- 
guished by  these  tests,. 


76  LABORATORY    MANUAL. 

Experiment  71.— Sulphuric  Dioxide. 

Materials  and  Apparatus. — Deflagration-spoon,  sulphur,  large  bottle, 
a  red  rose,  nitric  acid,  some  cotton,  sulphuric  acid. 

Light  a  small  lump  of  sulphur  in  the  deflagration-spoon  and  lower  it 
into  a  large  wide-mouthed  bottle  covered  by  a  card.  As  sulphur  burns 
what  gas  is  formed1?  Compare  the  valency  of  the  two  elements.  "What 
does  the  odor  of  the  gas  resemble  ?  Will  it  burn  ?  Try  a  lighted  splinter. 
Hold  a  red  rose  or  any  red  flower  in  the  mouth  of  the  bottle.  What 
effect  1  Fill  a  beaker  with  water  and  pour  a  little  sulphuric  acid  into  it. 
[NOTE. — Never  pour  water  into  sulphuric  acid,  but  pour  the 
acid  into  water.  Why  ?]  Immerse  the  rose  in  the  dilute  acid. 
Dry  and  warm  it.  Is  the  color  restored  1  Try  to  bleach  wet 
calico,  also  blue  litmus  paper.  What  effect?  Does  sulphur 
dioxide  bleach  as  chlorine  does,  by  combining  with  the  hydro- 
gen of  water,  setting  free  the  oxygen,  the  latter  burning  up 
the  coloring-matter  ?  Straw  hats  are  bleached  by  sulphur. 
Why  do  they  become  yellow  when  wet?  Why  do  flannels 
become  yellow  1  By  means  of  a  wire  or  glass  rod  lower  a  piece 
of  cotton  or  sponge  wet  with  nitric  acid  into  the  bottle  of  gas.  The  red 
fumes  are  the  result  of  reduction  or  the  taking  away  of  oxygen  from  the 
acid,  thus  :  2HN03  +S02  =  H2S04  +  2N02-  Give  the  name  of  each  as 
indicated. 

Does  this  suggest  what  might  be  used  in  making  sulphuric  acid  ? 

Experiment  72.— Phosphorus. 

Materials  and  Apparatus. — Phosphorus,  carbon  bisulphide,  beaker, 
filter-paper. 

Put  a  lump  of  phosphorus  as  large  as  a  grain  of  wheat  into  a  beaker 
and  pour  upon  it  immediately  a  c.  c.  of  carbon  bisulphide.  Does  it 
dissolve  ?  Pour  a  few  drops  of  the  solution  upon  a  filter-paper  and 
allow  it  to  evaporate.  The  paper  may  be  held  upon  the  palm  of  the 
hand  if  care  is  taken  not  to  hold  it  too  near  the  person,  so  that  the 
bursting  into  flame  will  not  set  fire  to  the  clothing.  Why  does  it  burn 
spontaneously  ?  Why  does  not  the  phosphorus  on  the  head  of  a  match 
burn  spontaneously  1  What  use  have  you  seen  made  of  phosphorus  in 
preparing  other  elements?  Why  is  phosphorus  kept  under  water? 
Why  does  it  not  burn  up  the  oxygen  contained  in  the  water  ? 

NOTE. — Teacher  perform  Experiment  59,  Eliot  and  Storer's  "Elementary  Chem- 
istry," page  96. 


78  LABORATORY    MANUAL. 

Experiment  73.— Arsenic, 

Materials  and  Apparatus. — A  hard  glass  tube,  some  charcoal,  and 
white  arsenic. 

Seal  up  one  end  of  a  piece  of  glass  tubing  four  inches  long.  Drop 
into  this  a  little  white  arsenic,  then  cover  with  some  small  lumps  of 
charcoal.  Heat  the  charcoal  red-hot  first,  then  heat  the  arsenic  below  it. 
What  is  the  effect  ? 

What  power  has  carbon  when  hot  with  reference  to  oxygen  in 
other  compounds  at  hand? 

What  is  the  metallic  substance  above  the  charcoal  ? 

Experiment  74.— Test  for  Arsenic. 

Materials  and  Apparatus. — Test-tube,  zinc,  sulphuric  acid,  sodium 
hydrate,  mercuric  chloride,  filter- paper. 

Put  some  grains  of  zinc  into  a  test-tube  and  cover  with  dilute  sul- 
phuric acid. 

Add  a  solution  of  white  arsenic,  made  by  boiling  the  powder  in 
sodium  hydrate  and  water. 

As  hydrogen  escapes  from  the  acid  by  the  action  of  zinc,  it  will 
combine  with  the  arsenic  and  escape  as  a  gas,  AsH3,  arseniureted 
hydrogen. 

Cover  the  mouth  of  the  test-tube  with  a  filter-paper  moistened  with 
mercuric  chloride,  and  in  a  few  moments  the  paper  will  be  stained 
yellow  to  orange  red.  (Characteristic  test.) 

Experiment  75,— Carbon  Dioxide. 

Place  a  short  piece  of  candle  on  your  desk ; 
light,  and  hold  a  dry,  clean,  wide-mouthed  bottle 
down  over  the  flame. 

What  do  you  see  formed  on  the  inner  wall  of 
the  bottle  1 

Kemove  bottle,  pour  in  a  small  quantity  of  lime- 
water,  and  shake. 

What   color    is  given   to  the   lime-water'? 

Fj    91  In  what  experiment  has  this  change  been  pro- 

duced before  1 

What  substance  is  formed  ? 
Write   the  equation. 

What  two  compounds  are  seen  to  be  the  products  of  combustion? 
Keview  Experiments  49  and  50. 


80 


LABORATORY   MANUAL. 


Experiment  76.— Preparation  of  C02  for  Use. 

Place  several  small  lumps  of 
marble  or  lime-stone  (calcium  car- 
bonate) in  generating  flask  used 
for  making  H.  Cover  with  water, 
and  pour  in  strong  HC1  until  gas 
is  given  off  abundantly.  Collect 
in  jars  over  water,  as  in  the  case  of 
hydrogen.  Collect  some  in  wide- 
mouthed  bottle  by  displacement 
of  air,  having  bottle  on  your  table 
with  mouth  up,  and  delivery-tube 
running  to  the  bottom. 

Complete  the  equation,  CaC03  4* 
of  this  gas  upon  lime-water.  Name 
lighted  candle  in  bottom  of  a  tumbler 


Fig.  22. 

2HC1  =  ?  Try  the  effect  of  some 
of  the  gas  1  Stand  short  bit  of 
and  pour  gas  upon  it  as 
you  would  water.  What  is 
the  effect  upon  the  flame  ? 
What  does  this  last  action 
show  as  to  weight  of  this 
gas  ?  Find  the  weight  of 
this  gas,  if  you  have  not 
done  so  already,  by  balanc- 
ing a  large  beaker  on  the 
scale-pan  and  filling  with 
CO  2  by  displacement  of  air. 
How  would  you  pour  H 
from  one  vessel  to  another 
containing  air  ?  In  what 
position  would  you  hold 
the  bottles  1 


Fig.  23.— WEIGHING  CARBON  DIOXIDE. 


Experiment  77.— Absorption  of  Carbon  Dioxide. 

Fill  a  bottle  half  full  of  C02  over  water  in  a  pneumatic  trough. 
Cork  the  bottle  under  water  and  shake.  Lower  mouth  of  bottle  into 
water  again,  remove  stopper,  and  note  the  rise  of  water  on  the  inside. 
To  what  is  this  due  ? 

What  is  "  soda-water "  ?  What  other  gas  which  you  have  studied 
was  dissolved  in  water  ? 


82  LABORATORY  MANUAL. 

Experiment  78.— Saturation  with  Carbon  Dioxide, 

Conduct  some  C02   from    delivery-tube  into  beaker  of  lime-water. 

Note  whitening  as  before.  What  substance  is  formed  ?  Is  it  soluble  in 
water  ? 

Continue  to  pass  the  gas  until  white  color  seems  to  disappear.  Can 
you  explain  1  What  does  water  contain  after  the  gas  has  passed  through 
it  1  What  is  the  effect  of  water  in  this  condition  upon  limestone  ? 

Heat  the  contents  of  the  beaker  and  see  if  the  white  substance  reap- 
pears. When  some  water  is  boiled  in  a  tea-kettle,  limestone  is  deposited 
upon  the  bottom  and  sides.  Explain.  What  gas  does  such  \vater  contain 
in  solution  before  boiling  ?  In  caves  and  limestone  regions  the  water,  in 
coming  to  the  surface,  deposits  limestone  in  the  form  of  icicles,  some 
hanging  down  from  above  and  others  rising  from  the  floor.  Can  you 
give  the  explanation  ? 

How  does  dough  rise  ?     (See  Experiments  16  and  90.) 

Experiment  79.— Candle  Flame. 

Examine  the  flame  of  a  candle.  Have  it  as  large  as  possible,  and 
protected  from  draughts  of  air.  Note  (1)  the  blue  cup-shaped  portion  at 
the  base,  (2)  the  dark,  non-luminous  portion  inside  surrounded  by  the  (3) 
luminous  cone,  and  (4)  outside  of  that  by  a  thin  layer  or  region  of  com- 
plete combustion.  Make  a  sketch  of  the  flame  in  your  note-book,  show- 
ing all  its  parts. 

NOTE. — The  spectrum  given  by  a  solid  or  liquid  in  a  state  of  incandescence  is  a 
continuous  band  of  color.  That  given  by  a  gas  in  a  burning  state  consists  of  bright 
lines  here  and  there  across  the  place  occupied  by  the  continuous  spectrum  when 
present. 

Examine  the  candle  flame  with  a  spectroscope.  Does  it  show  the 
light-giving  portion  of  the  flame  to  consist  of  a  solid  or  liquid,  on  the  one 
hand,  or  of  a  gas  1 

Hold  a  cold  iron  spoon  down  upon  the  flame  until  it  almost  touches 
the  wick.  What  collects  upon  the  spoon  ?  Next  hold  this  blackened 
surface  for  some  time  in  the  non-luminous  flame  of  the  Bunsen  burner, 
and  describe  what  takes  place.  What  is  your  conclusion  as  to  the  con- 
dition of  matter  in  the  luminous  portion  of  a  flame  1  At  which  portion 
of  the  flame  does  the  air  have  freest  access  to  the  combustible  material  ? 
Where  is  the  region  of  complete  combustion  ?  Is  this  luminous  or  non- 
luminous  ?  Press  down  upon  the  flame  a  piece  of  white  paper  or  card- 
board. Hold  it  there  for  an  instant  and  remove  quickly  to  avoid  taking 
fire.  What  does  this  show  in  regard  to  the  interior  of  the  flame  ?  Where 
was  the  paper  charred  1  Sketch  cross-section  in  your  notes. 


84 


LABORATORY   MANUAL. 


Fig.  24. 


Experiment  80.— The  Flame. 

(a)  Stand  a  lighted  candle  on  smooth  table, 
and  place  a  tall,  narrow  lamp-chimney  over  it, 
so  that  no  air  can  enter  from  below. 

Does  the  candle  continue  to  bum  ] 

(b)  Now  raise  edges  of  chimney  up  on  thin 
strips  of  wood  or  metal  so  that  air  may  enter. 
What  is  shown  by  this  experiment  ? 

Now,  with  candle  as  in  this  experiment  (b), 
cover  chimney  with  strips  of  metal  or  evaporat- 
ing dish.  Explain  what  happens. 


Experiment  81. — Kindling  Temperature. 

Lower  spoon  in  a  candle  flame  almost  to  the  wick,  and  collect  carbon 
upon  it.     Why  is  carbon  deposited,  and  why  does 
it  not  continue  to  burn  up  ?     What  is  "  kindling 
temperature  "  1 

Fold  a  piece  of  commercial,  note-paper  as  you 
do  your  filter-paper.  Pour  a  little  water  into  it, 
and  place  in  one  of  the  rings  of  your  ring-stand, 
and  boil  water  with  Bunsen  burner. 

What  is  temperature  of  boiling  water?  What 
must  be  the  temperature  of  the  paper  next  to  it  ? 
Is  this  temperature  as  high  as  the  kindling  temper- 
ature of  the  paper  1  Explain  the  purpose  of  using 
a  match,  pieces  of  paper,  and  wood,  in  this  order,  in  building  a  coal  fire. 

Experiment  82.— The  Bunsen  Burner. 

When  the  holes  are  closed  at  the  bottom  of  the  lamp,  light  the  gas 
and  examine   the   flame.     Compare   with   candle  flame.     Explain   the 

change  in  character  of  the 
flame  produced,  with  the 
orifices  at  the  bottom  of 
the  lamp  open.  Find  the 
hottest  part  of  this  flame 
by  means  of  a  platinum 
wire.  Lower  a  piece  of 
wire-gauze  down  upon  the 
flame,  and  notice  structure 
of  the  flame  in  cross-section 
by  seeing  where  the  wire- 
gauze  is  made  to  glow. 

Fig.  26. — BUNSEN  BURNER. 


86 


LABORATORY   MANUAL. 


Fig.  28.— DAVY'S  SAFETY - 
LAMP. 


Experiment  83.— Principle  of  Sir  Humphry 
Davy's  .Safety-lamp. 

Arrange  wire-gauze  on  ring-stand,  as    shown  in 
Fig.  27.      Turn  on  the  gas  and  light  it  above  the 
gauze.     Is  there  gas  below  the  gauze  ? 
Prove  it. 

Extinguish  flame  and  light  the  gas 
below  the  gauze.  Does  the  flame  pass 
through  it  1  Explain. 

Look  up  the  subject  of  "  Construction 
and  Use  of  Davy's  Safety-lamp,"  page  134 
of  Eliot  and  Storer's  "Elementary  Chemis- 
Fig.  27.  try,"  and  elsewhere. 

Experiment  84.— Action  of  Hot  Carbon  upon 
Oxygen  in  Combination. 

Mix  thoroughly  .25  gram  powdered  charcoal  and 
2.5  grams  copper  oxide,  and  place  in  a  hard  glass 
test-tube  provided  with  a  delivery-tube. 

Heat  persistently  over  Bunsen  flame,  and  con- 
duct gas  into  another  test-tube  containing  lime-water. 
What  is  the  gas  ?  After  gas  ceases  to  be  given  off,  examine  contents  of 
hard  glass  tube.  Is  it  copper  ?  By  weighing  tube  before  and  after  heat- 
ing, find  weight  of  copper  in  2.5  grams  of  copper  oxide. 


Experiment  85.— The  Blow-pipe. 

NOTE.  — Kead  carefully  Eliot  and  Storer's  ' '  Elementary 
Manual  of  Chemistry,"  pages  129  and  130. 

Place  a  small  quantity  PbO  in  a  hollow  made  in 
piece  of  charcoal.  Heat  in  the  "  reducing  flame  "  for 
some  time,  until  minute  drops  of  metal  are  seen  in 

place    of    the 

--si;  c       b       ^jtf-""^^— |^^^s     powder.    What 

is  the  metal  1 
Try  to  cut  it 
with  a  knife. 
Write  equation 
for  above  re- 
Fig.  29.— COMMON  BLOW-PIPES.  action. 

NOTE. — Experiments  84  and  85  illustrate  in  general  the 
method  of  securing  metals  from  their  ores,  as  metallic  iron 


from  iron  oxide. 


88  LABORATORY   MANUAL. 


Fig.  30. — USE  OF  BLOW-PIPE. 

Experiment  86.— The  Oxidizing  Flame, 

Place  a  small  bit  of  lead  on  charcoal  hollowed  out  as  in  preceding 
experiment,  and  heat  in  the  oxidizing  flame.  What  is  the  name  of  the 
coating  found  about  the  lead  on  the  charcoal  ?  What  substance  does  it 
resemble?  Write  the  equation.  Has  the  same  substance  been  produced 
before  ?  In  Avhat  experiment  1 

Experiment  87.— The  Borax  Bead. 

Follow  the  directions  for  making  a  borax  bead  given  in  Eliot  and 
Storer's  "Elementary  Chemistry,"  pages  217  and  218.  Touch  the  bead 
to  a  small  speck  of  black  oxide  of  manganese,  and  heat  in  the  oxidizing 
flame.  Pulverize  this  bead  between  heavy  pieces  of  iron.  Be  careful 
not  to  bend  the  platinum  wire.  Clean  the  wire  and  make  a  new  bead. 
Dip  this  in  cobaltous  nitrate  solution,  or  in  the  salt  itself,  and  heat  as 
before.  Note  color  when  hot  and  when  cold. 

Borax-bead  tests  for  other  metals  will  be  referred  to  in  Experiments 
on  Metals,  further  along  in  this  book. 

NOTE. — Borax  dissolves  the  oxides  of  various  metals;  hence  its  use  in  prepar- 
ing a  clean  surface  in  soldering. 

Experiment  88.— Marsh  Gas. 

(For  the  Teacher.} — Perform  the  experiment  as  directed  on  page  138 
Eliot  and  Storer's  "  Elementary  Chemistry." 

Why  does  it  explode  when  mixed  with  air  and  ignited  ?  In  what 
proportions  do  they  explode  best  ?  What  must  be  formed  when  CH4 
burns  ?  How  can  you  prove  that  water  is  one  of  the  products  1  C02  ? 
What  proof  that  carbon  is  left  when  chlorine  and  CH4  explode  1  Why 
do  miners  call  this  gas  fire-damp  ?  Why  does  Davy's  safety-lamp  pre- 
vent an  explosion  ?  Under  what  conditions  will  an  explosion  occur  ? 
This  gas  forms  a  large  part  of  illuminating  gas.  Try  to  collect  the 
gas  from  a  pond.  Describe  the  whole  experiment. 


90  LABORATORY    MANUAL. 

Experiment  89,— Uluminating  Gas. 

Materials  and  Apparatus. — Two  clay  pipes,  plaster  of  Paris,  soft 
coal,  sawdust. 

Fill  a  clay  pipe  half  full  of  small  pieces  of  soft  coal.  Stop  the 
mouth  of  the  bowl  with  a  paste  made  of  plaster  of  Paris  and  water.  As 
soon  as  the  plaster  hardens  support  the  pipe  so  that  it  can  be  strongly 
heated.  At  first,  steam  will  issue  from  the  stem  of  the  pipe,  and  finally, 
gas.  Ignite  from  time  to  time  until  the  gas  becomes  pure  enough  to 
burn  of  itself. 

After  all  the  gas  is  driven  off  let  it  cool,  then  cut  out  the  plaster  and 
examine  the  contents  of  the  pipe.  What  is  the  black  mass  ?  Do  the 
same  with  some  sawdust  or  fine  chips.  Do  they  yield  a  gas  that  will 
burn  1  What  is  left  after  the  burning  or  heating  is  finished  1  How 
does  coke  differ  from  charcoal  ?  Why  did  not  the  wood  burn  to  ashes  1 
How  is  charcoal  prepared  for  commerce  ?  Some  of  the  impurities  formed 
when  coal  is  thus  heated  to  make  illuminating  gas  are  C02  and  NH3. 
How  could  these  two  impurities  be  removed  1  Does  CH^  from  the 
pipe-stem  burn  with  a  clear  or  smoky  flame  1  Is  carbon  wasting  ?  How 
can  this  be  remedied  1  Visit  some  gas-works  and  see  how  gas  is  made 
on  a  large  scale. 

Experiment  90. —Alcoholic  Fermentation. 

Dissolve  30  grains  of  grape-sugar  (C6H1206)  or  cane-sugar 
(CjgH^On)  in  400  c.  c.  of  water.  Put  the  solution  into  a  flask  pro- 
vided with  a  stopper  and  delivery-tube  to  conduct  the  gas  generated 
into  a  bottle  of  lime-water  (CaH.,02). 

Add  to  the  solution  of  sugar  a  fourth  of  a  cake  of  compressed  yeast 
or  a  tablespoonful  of  baker's  yeast.  Insert  the  stopper  and  place  the 
apparatus  in  a  warm  place  for  from  twenty-four  to  forty-eight  hours.  Does 
the  lime-water  indicate  that  C02  has  been  generated  ?  Taste  the  solu- 
tion. Is  it  sweet?  The  chemical  change  may  be  indicated  by  the 
equation,  C6H1206  =  2C2H60  +  2C02.  That  is,  by  the  action  of 
yeast,  grape-sugar  splits  up  into  two  molecules  each  of  alcohol  and  car- 
bon dioxide. 

Or  if  cane-sugar  was  used  it  changes  first  to  grape-sugar,  thus  : 
C12H22011  +  H,0  =  2C6H1206. 

From  what  source  is  alcohol  obtained  1  Since  the  boiling  point  of 
alcohol  is  20°  lower  than  that  of  water,  can  you  suggest  a  method  of 
separating  alcohol  from  water  1 

Save  the  dilute  alcohol  you  have  made  for  the  next  experiment. 


92  LAB  OR  A  TORY    MANUA  L. 

Experiment  91.— Vinegar. 

Allow  the  dilute  alcohol  made  in  the  last  experiment  to  stand  sev- 
eral days  exposed  to  the  air.  Finally,  it  will  acquire  the  characteristic 
taste  of  vinegar.  The  reaction  is:  C2H60  +  02  =  HC2H302  +  H20. 
From  what  source  is  vinegar  (acetic  acid)  obtained  ?  Sometimes  water 
is  made  sour  by  a  small  quantity  of  sulphuric  acid.  How  could  you 
detect  it?  (Bee  Sulphuric  Acid,  page  72.)  What  is  yeast?  Is  the 
change  from  alcohol  to  vinegar  due  to  yeast  1  How  can  you  prove  it  ? 

Experiment  92.— Fractional  Distillation. 

(For  the  Teacher.) — Collect  part  of  the  dilute  alcohol  from  the  class 
or  prepare  some  in  a  similar  way.  Arrange  apparatus  as  shown  in  Eliot 
and  Storer's  "  Elementary  Chemistry,"  page  146.  Distill  several  times  to 
get  sufficient  strength  of  alcohol  to  burn.  Put  a  few  drops  of  alcohol 
upon  the  hand.  Why  is  it  cold1?  Drop  some  camphor-gum  into  a  few 
c.  c.  of  alcohol.  Does  it  dissolve  ?  Add  water  and  it  will  be  precipitated. 
Pour  strong  alcohol  upon  white  of  egg.  What  effect  1  Put  some  green 
leaves  into  a  bottle  and  cover  with  alcohol.  Allow  to  stand  over  night. 
Is  the  green  coloring  matter  (chlorophyl)  extracted?  Mix  together 
two  volumes  of  concentrated  alcohol  and  two  of  water.  Does  it  make 
four  volumes  of  the  mixture  1  How  do  you  account  for  it  ?  How  does 
alcohol  preserve  museum  specimens  ? 

Experiment  93.— Test  for  Sugars, 

Dissolve  rock-candy  or  granulated  sugar  in  water.  Take  5  c.  c.  of 
the  solution  in  a  test-tube.  In  another  test-tube  take  5  c.  c.  of 
grape-sugar  (a  raisin  soaked  in  water  will  give  you  grape-sugar).  To 
each  of  these  test-tubes  add  two  2  c.  c.  of*Fehling's  solution  of 
copper  sulphate.  Heat  gently  over  the  burner.  Which  one  gives 
a  yellow  to  brown  precipitate  of  copper  hydrate?  Add  a  few  drops 
of  H2S04  to  the  granulated  sugar  (cane-sugar) ;  boil.  Does  it  give  the 
yellowish  precipitate  ?  This  is  a  delicate  test  for  grape-sugar.  Boiling 
converts  the  yellow  hydrate  into  a  dark -brown  sub-oxide.  What  must  be 
the  action  of  Ho  SO 4  upon  cane-sugar  (Cj  2H2  gOj  j)  jn  order  to  change  it  to 
SCgHj  206  ?  Apply  the  copper  test  to  the  following  to  determine  whether 
the  sugar  is  cane-  or  grape-sugar  :  Honey,  peach,  or  any  other  fruit-juice, 
molasses,  various  samples  of  candy.  Try  prolonged  boiling,  and  see  if  that 
will  change  cane-  to  grape-sugar.  Put  yeast  into  a  solution  of  cane-sugar 
and  allow  to  stand  twenty-four  hours.  Has  it  changed  to  grape-sugar  ? 

NOTE  TO  TEACHER.  —  Try  Experiment  134,  page  180,  Eliot  and  Storer's 
"  Elementary  Chemistry,"  and  test  for  sugar  with  Fehling's  solution.  Try  to  con- 
vert cloth  or  wood  into  sugar  in  the  same  way. 

*  Page  140. 


94  LABORATORY    MANUAL. 

Experiment  94.— Digestive  Ferments,  Albuminous. 

NOTE. — The  digestive  ferments,  pepsin,  of  the  gastric  juice,  pancreatin,  of  the 
pancreas,  and  amylopsin,  one  of  the  three  ferments  of  pancreatin,  may  be  obtained 
from  the  druggist.  They  are  prepared  in  powder  and  may  be  kept  any  length  of 
time.  By  the  following  experiments  try  to  determine  which  of  the  ferments  acts 
upon  each  of  the  classes  of  food  ;  which  acts  most  rapidly  ;  what  effect  heat  and 
cold  have  upon  digestion;  etc.  The  experiments  may  be  performed  by  each  pupil,  or 
the  teacher  may  perform  the  experiments  before  the  class,  the  pupils  taking  notes 
as  directed. 

Boil  an  average-sized  egg  ten  or  fifteen  minutes.  Separate  the  yolk 
from  the  white  and  rub  each  separately  through  a  thirty-mesh  sieve. 
Divide  each  into  three  portions  by  weight  and  put  them  into  beakers 
or  into  two-ounce  flat  bottles.  Number  the  bottles  containing  the 
white  1,  2,  and  3  respectively,  and  those  containing  the  yolk,  4,  5,  and  6. 
Add  to  each  of  these  25  c.  c.  of  one-per-cent.  HC1.  Add  to  Nos.  1  and 
4  one  centigram  or  more  each  of  pepsin  ;  to  Nos.  2  and  5  one  centigram 
each  of  pancreatin ;  and  to  Nos.  3  and  6  one  centigram  each  of  amylopsin. 
Shake  thoroughly  and  set  in  a  pan  of  water  kept  at  a  temperature  of 
120°  to  130°  F.  Agitate  every  five  minutes.  In  which  does  the  albumin 
dissolve?  In  which  does  the  oily  yolk  dissolve?  When  you  are  satis- 
fied as  to  which  are  not  changed  by  the  digestive  ferment,  neutralize  these 
fluids  with  sodium  carbonate  and  replace  in  the  water-bath.  After  suffi- 
cient time  has  elapsed  examine  again.  In  which  cases  now  does  the 
food  dissolve  in  the  neutralized  fluid  ?  Is  the  gastric  juice  acid  or  alka- 
line ?  Are  the  intestinal  fluids  the  same  ?  What  class  of  food  is  digested 
in  the  stomach  ?  What  is  the  temperature  of  the  human  body  ? 

Experiment  95.— Effect  of  Temperature  upon  Digestion, 

Prepare  an  egg  and  apparatus  as  in  the  last  experiment,  except  that 
only  the  white,  the  HC1,  and  pepsin  will  be  required.  Place  No.  1  in 
ice-water,  No.  2  in  water  at  100°  to  105°  F.,  and  No.  3  in  water  at  130°  F. 
In  which  does  the  albumin  dissolve  most  rapidly  ?  In  which  most  slowly  ? 
What  effect  upon  digestion  does  hot  tea,  coffee,  etc.,  have,  at  least  so  far 
as  temperature  is  concerned  ?  How  would  hot  food  compare  with  cold 
food  ?  How  would  the  drinking  of  ice-water  affect  digestion  for  a  time  ? 

Experiment  96.  —Effect  of  Mastication  upon  Albuminous  Food. 

Prepare  an  egg  as  before,  except  that  one  third  of  it  should  be  in 
a  solid  chunk ;  a  second,  cut  into  small  pieces ;  a  third,  passed  through 
the  finest  sieve.  Be  careful  to  use  the  same  quantity  of  albumin  and 
pepsin  and  acid  for  each,  and  keep  them  all  at  the  same  temperature 
in  a  water-bath.  In  which  does  digestion  take  place  most  rapidly? 
In  which  most  slowly  ?  Why  ?  Which  would  resemble  rapid  eating  ? 


96  LABORATORY    MANUAL. 


Experiment  97.— Digestion  of  Starches. 

Prepare*  iodine  solution  and  Fehling's  fluid.  Make  starch-paste  by 
1  oiling  about  30  grams  of  laundry  starch  in  about  500  c.  c.  of  water. 

Boil  until  it  becomes  translucent,  stirring  thoroughly.  Put  about 
100  c.  c.  of  this  paste  into  each  of  five  bottles,  and  stand  four  of  them 
in  the  water-bath,  to  be  kept  at  a  temperature  of  105°  to  107°  F. 

Number  the  bottles  1,2,  3,  4,  and  5  respectively.  Place  No.  5  in  ice- 
water.  Weigh  out  carefully  one  decigram  of  pepsin  for  No.  1 ,  one  decigram 
of  pancreatin  for  No.  2,  one  decigram  of  amylopsin  each  for  3  and  5. 

Put  these  weighed  quantities  of  the  ferments  into  their  respective 
bottles,  and  shake  thoroughly.  Notice  that  4  is  only  for  comparison. 

How  does  the  ferment  in  any  case  affect  the  starch  mass  ?  In  which 
does  it  liquefy  most?  Take  five  ordinary  test-tubes  and  fill  with  water. 
Set  them  in  a  row,  or  number  with  a  gummed  label.  Drop  one  small 
drop  of  iodine  solution  into  each  by  means  of  a  pipette,  and  shake  to 
mix. 

Now,  when  the  ferments  have  been  at  work  for  five  minutes,  take  a 
clean  pipette  and  drop  one  drop  into  iodine  solution  No.  1.  Rinse  the 
pipette  and  do  the  same  with  each  number  successively.  Then  shake 
the  iodine  solution. 

No.  4  will  show  what  pure  starch  and  iodine  will  give.  How  do  the 
others  compare  ?  In  which  do  you  find  the  least  change  of  color  by  the 
iodine  ?  Allow  to  stand  five  minutes  more,  and  repeat  the  same,  having 
replenished  your  test-tubes  with  iodine  solution.  What  is  the  effect  of 
temperature  as  shown  by  the  ice-water1? 

After  you  have  settled  the  question  as  to  which  of  these  three 
ferments  is  most  active  in  digesting  starch,  test  the  product  of  digestion 
to  determine  what  the  starch  has  been  changed  into.  Take  about  5  c.  c. 
of  Fehling's  solution,  freshly  made.  Boil  it  in  a  test-tube,  to  be  sure 
it  is  good.  It  should  not  change  color. 

Now  add  to  it  a  few  drops  of  digested  starch,  and  boil.  What 
does  this  show  it  to  be  ?  Test  each  of  the  bottles  in  the  same  way. 

In  which  is  there  no,  or  very  little,  change,  and  in  which  most  ?  Does 
this  agree  with  the  starch  test  ?  Where  in  the  alimentary  canal  is  this 
change  made? 

Chew  a  cracker  for  a  few  minutes.  Does  it  become  sweeter  ?  What 
does  this  show  ?  Does  the  result  with  regard  to  temperature  agree  with 
the  last  ? 

Cane-sugar  is  converted  into  grape-sugar  before  being  absorbed,  but  it 
is  not  known  by  what  means  it  is  changed. 

*  Page  140. 


98  LABORATORY   MANUAL. 


Experiment  98.— Effect  of  Acids  and  Alkalies  upon  Starch  Digestion, 

Take  four  bottles  of  starch,  prepared  as  above,  and  make  1  and  2  acid 
with  one-per-cent.  HC1,  and  3  and  4  alkaline  by  NasC03.  Number  and  add 
the  same  portion  as  above  of  pepsin  to  1  and  3,  and  amylopsin  to  2  and  4. 
Keep  all  at  the  temperature  of  105°  F.,  and  test  as  before.  What  do 
you  learn  as  to  the  effects  of  acids  and  alkalies  upon  these  ferments  in 
their  action  toward  starches  ? 


Experiment  99.— Digestion  of  Milk. 

Prepare  the  water-bath,  and  heat  the  water  to  105°  F.  Put  25  c.  c. 
of  fresh  milk  in  each  of  the  three  test-tubes,  and  leave  them  in  the  bath 
long  enough  to  become  heated  to  that  temperature.  To  3  c.  c.  of  water 
add  3  decigrams  of  sodium  carbonate  and  1  decigram  of  pancreatin,  and 
mix  thoroughly. 

Put  this  into  No.  1,  shake,  and  replace  in  the  bath.  Add  to  No.  2 
a  like  preparation  of  pepsin  and  sodium  carbonate,  and  to  No.  3,  3  c.  c. 
of  one-per-cent.  HC1  and  1  decigram  of  pepsin.  We  thus  have  the  acid 
ferment  of  the  stomach,  No.  3,  the  same  neutralized,  No  2,  and  the  natural 
alkaline  ferment  of  the  intestine. 

How  "do  they  appear  in  the  course  of  five  minutes? 

Which  one  has  formed  a  clot  or  curd  ? 

Which  one  seems  to  dissolve  up  the  milk  1 

To  test  it,  take  5  c.  c.  of  pure  milk,  and  add  a  drop  of  dilute  HN03, 
and  it  forms  a  clot.  Do  the  same  with  No.  1.  Is  the  milk  changed  1 
Is  there  a  clot  or  a  precipitate  formed  by  HN03  1 

Allow  to  stand  for  an  hour.  How  have  they  progressed  now  ?  Test 
for  sugar  by  Fehling's  fluid.  In  which  has  tho  milk  been  prepared 
most  for  absorption?  Does  acid  or  alkaline  pepsin  digest  milk  most 
rapidly  ? 

If  milk  forms  such  a  hard  clot  in  the  stomach,  what  could  be  done  to 
make  it  digest  more  quickly  ?  Would  lime-water  taken  with  it  prevent 
this  ?  Human  milk  does  not  form  such  a  clot  as  this  in  an  infant's 
stomach,  but  remains  in  small  threads. 

By  the  addition  of  pancreatin  to  the  cow's  milk  a  few  minutes  before 
taking,  it  may  be  kept  from  forming  hard  curds.  Is  the  coagulation  of 
milk  due  to  the  ferment  or  to  the  HC1  of  the  stomach?  Try  milk  and 
HC1.  Which  of  the  two  ferments  digests  milk  best  ?  What  is  cheese  1 
How  is  rennet  obtained  1 


100  LABORATORY  MANUAL. 

Classifications  of  Elements  and  Compounds. — There  tire  about  twenty-three  ele- 
ments that  act  as  bases  to  form  the  most  common  salts,  and  since  by  far  the  greatest 
number  of  compounds  are  salts,  it  is  important  to  know  the  characteristics  of  each 
base,  so  that  we  may  determine  the  composition  of  an  unknown  compound.  If  we 
can  perform  an  experiment  by  means  of  some  reagent  by  which  the  character  of  the 
base  and  the  acid  radical  are  brought  out,  we  may  then  apply  the  laws  of  combina- 
tion, and  say  of  what  the  salt  is  composed. 

These  characteristics  are  very  largely  the  difference  in  solubility,  color,  odor,  etc. 

In  all  the  following  experiments  notice  carefully  these  differences  and  resem- 
blances, and  make  full  notes,  so  that  you  may  refer  to  them  again  and  gather  up 
groups  of  facts  as  directed.  Write  out  all  equations  in  full,  referring  back  to  the 
Table  of  Elements,  page  28,  if  you  do  not  know  the  atomicity  of  an  element,  and  to 
the  Table  of  Basicity  of  Acids,  page  34,  if  you  do  not  know  the  combining  power  of 
an  acid.  Remember  that  all  that  you  take  must  be  in  the  first  member  of  the  equa- 
tion, and  all  you  get  in  the  second  member,  and  that  the  second  member  cannot 
contain  an  element  or  a  different  quantity  of  the  element  that  did  not  appear  in 
the  first  member.  In  other  words,  you  cannot  make  or  destroy  an  atom  of  matter. 


Experiment  100.— Sulphides. 

NOTE  TO  TEACHER. — Prepare  HSS  upon  a  large  scale,  and  keep  the  generating 
flask  in  a  gas-hood  or  window  where  the  escaping  gas  will  be  carried  off.  Provide 
the  flask  with  a  delivery-tube  to  conduct  the  gas  to  the  bottom  of  a  test-tube. 
Keep  a  beaker  of  water  beside  the  flask,  so  that  the  delivery-tube  may  be  washed 
after  each  experiment,  to  prevent  mistakes. 

Make  a  test-tube  full  of  the  solution  of  each  base  to  be  tested.  First 
try  cold  water ;  then,  if  no  solution  is  formed,  boil.  If  none  still,  try 
diluted  HC1  and  if  heat  will  not  bring  about  the  desired  end,  add  HN03. 
If  it  still  refuses  to  dissolve,  heat  to  fusion  on  charcoal  with  Na2C03,  and 
then  it  will  dissolve  in  acids. 

We  wish  now  to  see  what  sulphides  we  can  form  by  precipitation 
from  a  solution  of  each  of  the  twenty-three  bases.  Some  of  them  form 
precipitates  with  H2S  from  an  acid  solution,  and  many  will  not.  Of  those 
that  will  not  thus  precipitate  sulphides  a  few  will  if  the  solution  is  neutral 
or  alkaline,  and  since  H2S  is  an  acid,  we  must  use  a  soluble  salt  of  that 
acid,  viz.,  (NH4)2S. 

To  determine  these  sulphides,  divide  your  test-tube  of  the  solution 
into  two  portions,  and  if  the  salt  was  dissolved  in  water  it  is  probably 
neutral,  since  most  salts  are  neutral.  Add  to  one  of  these  portions  of  the 
solution  a  few  drops  of  the  acid  of  the  same  kind  as  indicated  in  the 
radical  of  the  salt  to  acidify  it.  Thus,  if  the  salt  is  CuS04,  acidify  by  a 
few  drops  of  H2S04.  If  NaCl,  acidify  with  HC1,  etc.  If  the  solution  is 
already  acid,  neutralize  one  portion  with  NH4HO.  Complete  the  dia- 
gram upon  the  note  page,  and  try  each  of  the  twenty-three  common  bases, 
using  the  element  or  any  salt  of  it. 

For  a  reason  to  be  given  later,  try  to  dissolve  each  of  the  sulphides 
obtained  with  H2S  from  an  acid  solution,  with  (NH4)2S.  The  precipi- 


LABORATORY  MANUAL. 


101 


tates  obtained  with  aluminum  and  chromium  will  be  hydrates  instead  of 
sulphides  ;  but  this  fact  may  be  disregarded  here.  Try  each  of  the  bases 
as  follows  : 


SOLVENT. 

SALT  OR 
METAL. 

0 
H 

5 

O 
0 

Neut. 
Ac. 
Ac. 
Neut, 
Neut. 
A. 

REAGENT. 

PRECIPITATE 
AND  COLOR. 

EFFECT  OF  (NH1)2S. 

REMARK. 

H.O 
H20 
HC1 
HC1 
H,0 
H20 

CuSO, 
CuSO< 
SbCl3 
SbCl3 
FeSO, 
FeSO* 

(NH,),S 
H,S 

(NH~,)2S 

(NHJ,S 

Black  ppt. 

Orange  " 

Black     " 

Insoluble  in  (NHJ.S 
Soluble  in  (NHJ2S 

1 

102  LABORATORY   MANUAL. 

Experiment  101.— Chlorides. 

Prepare  a  table  similar  to  that  for  sulphides,  and  try  each  of  the  bases 
to  learn  the  same  facts  as  to  what  bases  form  chlorides  in  an  insoluble 
condition,  or  a  precipitate. 

You  will  notice  that  although  NH4HO  and  HC1  were  united  as  in 
Experiment  25  to  form  a  chloride,  yet  as  the  NaCl  is  soluble  in 
water,  no  precipitate  appeared.  By  evaporating  to  dryness  the  salt  was 
obtained. 

Now,  when  you  put  two  solutions  together  there  may  be  a  trading  of 
elements,  and  the  new  salt  formed  being  soluble  either  in  the  water  of 
the  solution  or  the  acid  present,  does  not  appear  as  a  precipitate. 

A  chemical  change  must  be  detected  by  some  other  means  than  by 
the  formation  of  a  precipitate. 

Again,  there  may  be  no  chemical  cjiange,  as  two  substances  may  not 
act  upon  each  other  at  all. 

Usually  the  evolution  of  a  gas  or  liberating  of  heat  will  show  that  a 
chemical  change  is  taking  place. 

Of  course  you  will  not  expect  to  precipitate  a  chloride,  bv  HC1  from 
a  chloride,  so  you  must  use  some  other  acid,  as  HN03,  to  form  a  solution 
if  it  cannot  be  done  by  water,  and  some  other  than  HC1  must  be  used 
for  acidifying. 

Do  not  be  discouraged  if  the  number  of  precipitates  is  small,  for 
there  are  but  few  chlorides  that  are  insoluble  in  water. 

Experiment  102.— Carbonates. 

Make  solutions  as  directed,  except  that  it  is  not  necessary  to  make 
acid  solutions,  since  the  reagent  by  which  the  carbonates  are  obtained,  is  a 
strong  alkali,  and  would  therefore  neutralize  the  solution. 

Since  H3C03  (carbonic  acid)  cannot  be  made,  we  must  obtain  car- 
bonates from  a  solution  of  a  salt  of  the  acid,  as  (NH4)2CO;r 

As  a  matter  of  economy,  it  is  best  to  neutralize  acid  solutions  before 
adding  ammonium  carbonate,  and  thus  save  a  large  quantity  of  the 
more  expensive  carbonate. 

Add  HN4HO  to  neutralize,  and  then  (HN4)2C03. 

Make  a  table  as  before.  Try  to  dissolve  the  few  precipitates  formed 
in  NH4C1. 

It  is  well  to  warm  the  solution  after  adding  the  ammonium  carbonate 
if  a  precipitate  does  not  form  readily. 

The  precipitates  obtained  will  all  be  carbonates. 


LABORATORY    MANUAL. 


103 


SOLVENT. 


SALT  OR 
METAL. 


REAGENT. 


PRECIPITATE 
AND  COLOR. 


EFFECT  OF 


REMARK. 


PRECIPITATES  OBTAINED  BY  (NH4)2COj 


SOLVENT. 

SALT  OR 
METAL. 

PRECIPITATE  AND 
COLOR. 

EFFECT  OF  NH«( 
UPON  PRECIPITAT 

HO 
H20 

MgSO< 
AgN03 

Milky-white  ppt. 
No  precipitate. 

Ppt.  dissolves. 

REMARK. 


Formed  slowly  only  af- 
ter heating. 


104  LABORATORY   MANUAL. 

Experiment  103,— Blow-pipe  Tests.    Metals, 

By  means  of  a  blow-pipe,  a  piece  of  charcoal,  and  some  dry  sodium 
carbonate,  and  a  gas-lamp  or  alcohol-burner,  determine  which  of  the  sub- 
stances indicated  give  metallic  globules,  which  form  a  coating  upon  the 
charcoal,  which  are  infusible,  which  give  off  vapors  while  heated,  what 
the  vapor  smells  like,  which  volatilize,  and  in  fact  all  the  changes  you 
notice. 

Salts  of  Ag,  Pb,  Hg,  As,  Sn,  Sb,  Cu,  Ni,  Fe,  Mg,  Zn. 

Experiment  104, — Blow-pipe  Test.    Salts. 

In  a  similar  way  to  the  last  experiment  determine  which  of  the 
salts  named  deflagrate  and  burn  the  charcoal,  which  fluff  up  and  form  a 
charred  mass,  which  give  off  the  odor  of  sulphur,  which  give  off  water,  as 
shown  by  fusing  and  then  becoming  dry,  etc.  ]STo  sodium  carbonate 
need  be  used  in  this  test ;  simply  the  charcoal,  plow-pipe,  and  lamp. 
Try  chlorates,  nitrates,  sulphates,  oxalates,  tartrates,  sulphides. 

Experiment  105.— Borax-bead  Test. 

Make  a  small  hook  on  the  platinum  wire  the  size  of  the  letter  o  in 
the  word  borax  in  the  head-line,  and  heat  it ;  thrust  it  into  the  borax- 
bottle  and  heat  till  you  have  a  clear  bead. 

Now  take  a  small  quantity  of  the  dry  powder  and  touch  the  hot 
bead  to  it.  Heat  it  again  until  it  becomes  clear.  Note  the  color. 
Make  a  new  bead  for  each.  Strike  the  bead  a  sharp  blow,  and  it  will 
crumble  out  of  the  loop. 

Which  of  the  salts  named  below  give  some  decided  color  to  the 
bead]  What  is  the  color  of  each?  FeS04,  Ni(N03)2,  CuS04,  NaCl, 
CO(N03)2,  Mn02,  K2Cr2Or 

Experiment  106 —Flame  Test, 

Pulverize  the  salts  to  a  fine  powder.  Moisten  the  platinum  wire  in 
HC1,  dip  it  into  the  powder,  and  burn  in  the  flame  of  the  lamp. 

Clean  the  wire  thoroughly  after  each  test,  either  by  burning  it  until 
it  gives  no  color  or  by  rubbing  it  with  emery  and  then  burning  it  a 
little. 

Note  the  color  of  the  flame  of  each  of  the  following  : 

CuS04,  BaCl2,  Sr(N03)2,  CaCl2,  KC1,  Na2C03. 

Now  try  other  salts  of  these  metals,  and  see  if  the  color  is  due  to 
the  base  or  to  the  acid  radical. 


106  LABORATORY   MANUAL. 


GROUPS    OF    BASES. 

If  the  work  has  been  carefully  and  properly  done,  you  should  have 
found  that  but  three  bases  form  chlorides  as  insoluble  precipitates  in 
water,  viz.,  Pb,  Ag,  and  Hg. 

We  may  use  this  means  of  separating  these  three  bases  from  the 
others,  and  for  convenience  we  will  call  this  the  First  Group,  or  the 
HC1  Group. 

The  acid  solution  of  bases  should  have  given  precipitates  with  As, 
Sb,  Sn,  B,  Cu,  Cd,  as  well  as  the  First  Group. 

Let  us  discard  the  three  bases  already  distinguished  from  the  rest  by 
their  chlorides,  and  we  have  a  Second  Group,  the  H2S  Group.  Three 
of  these  dissolved  upon  the  addition  of  (NH4)2S,  viz.,  As,  Sb,  arid  Sn, 
and  the  other  three  did  not. 

We  may  designate  the  soluble  sulphides  as  Division  A  of  the  H2S 
Group,  and  the  others  as  Division  1>. 

From  the  neutralized  solution  and  (NH4).,S,  were  formed  the  fol- 
lowing :  Fe,  Co,  Ni,  Zn,  and  Mn,  sulphides,  and  by  the  action  of  the 
reagent  Al  and  Cr  were  precipitated  as  hydrates. 

Classing  all  these  together,  we  have  a  Third  Group,  the  (NH4)2S 
Group. 

Ammonium  carbonate  should  give  precipitates  of  four  carbonates, 
Ba,  Sr,  Ca,  and  Mg,  the  latter  only  being  soluble  in  NH4C1. 

Call  these  four  the  Fourth  Group,  or  (NH4)2C03  Group. 

Of  the  twenty-three  common  bases  there  are  left,  the  two  elements 
K  and  Na,  and  the  radical  NH4,  which  forms  a  common  base. 

These  three  may  be  shown  to  be  closely  related,  and  are  therefore 
classed  as  the  Fifth  Group,  or  Alkaline  Group. 

Their  salts,  so  far  as  known,  are  all  readily  soluble  in  water  (see 
Table  I.,  page  107).  They  may  be  put  to  similar  uses,  as,  for  example, 
bicarbonate  of  soda,  or  bakers'  soda,  bicarbonate  of  potash,  or  saleratus, 
and  bicarbonate  of  ammonium  are  all  used  to  raise  biscuit. 

The  chlorides  of  each  differ  in  physical  properties,  chiefly  as  regards 
permanence  in  exposure  to  air  and  moisture.  The  hydrates  may  be 
used  for  the  same  purposes  in  softening  water,  neutralizing  acids,  etc. 

Before  Sir  Humphry  Davy  isolated  the  elements  K  and  Na,  they 
were  called  the  vegetable  alkali,  K,  the  mineral  alkali,  Na,  and  the 
volatile  alkali,  ammonium.  (For  the  analysis  of  an  unknown  base,  see 
Key,  page  112.) 

These  groups  are  only  arbitrary  divisions,  based  upon  the  resemblance 
with  reference  to  one  or  two  reagents,  called  group  reagents. 


LABORATORY    MANUAL. 


107 


•sapiqdpig 


ps  pt  pt 


*H  -*l  &  •*!  *4  *H  •*!    n  ^  <*<  *1  •*{•*<  ^  <*£•<*&• 


•sapixo     ^J 


^ 

•< 

§ 

p 

X 

C 


•sapipoi 


<j  ^ 


pc  ^  P* 


•sapiutojg; 


^  ^  <^  ^ 


^<<^^^^^^<-^<1^ 


108  LABORATORY   MANUAL. 

FIRST   GROUP   METALS. 
Pb,  Ag,  Hg'. 

SEPARATION    AND     IDENTIFICATION     OF    METALS     OF    THE    FIRST 

GROUP. 

Experiment  107. 

(a)  Measure  out  5  c.  c.  Pb(N03)2  solution  in  your  graduated  test-tube 
and  add  a  few  drops  of  HCL     What  salt  is  formed  ? 

(b)  Filter  and  wash  the  precipitate  with  a  little  cold  water. 

(c)  Pierce  the  bottom  of  the  filter-paper  with  your  platinum  wire, 
and  wash  a  quantity  of  the  precipitate  through  into  a  clean  test-tube. 

(d)  Boil  the  precipitate  in  water,  filter,  and  test  the  filtrate  by  adding 
a  few  drops  of  K2Cr04,  and  chrome  yellow  will  be  precipitated. 

Experiment  108. 

Repeat  the  above,  using  AgN03  solution  in  place  of  Pb(N03)2 
solution.  Is  PbCl2  dissolved  by  hot  water  ?  Is  AgCl  dissolved  in  hot 
water  1  Thus  we  have  a  means  of  separating  lead  and  silver  in  solutions 
of  their  salts. 

Experiment  109. 

(a)  Make  a  small  quantity  of  mercurous  chloride  and  try  to  dissolve 
it  in  hot  water.     Does  it  dissolve? 

(b)  Try  to  dissolve  a  small  quantity  of  AgCl  in  NH4OH  by  placing 
in  test-tube  and  heating. 

Try  same  experiment  with  Hg2Cl2.     Does  it  dissolve  in  NH4OH1 
Thus  we  have  a  method  of  separating  Pb,  Ag,  and  Hg  (cms)  in  solu- 
tions of  their  salts. 

Experiment  110. 

Take  5  c.  c.  solution  containing  all  three  salts,  Pb(N03)2,  AgN03,  and 
HgN03,  and  add  HC1.  Filter,  wash  precipitate  thoroughly  in  cold 
water,  and  throw  away  the  filtrate. 

(a)  Add  10  or  15  c.  c.  boiling  water  to  the  precipitate  on  filter-paper, 
and  test  filtrate  with  few  drops  of  K2Cr04.     Which  salt  is  dissolved  by 
the  hot  water  and   precipitated  upon  addition  of  K2Cr041     Write  the 
equation  for  the  last  reaction. 

(b)  To  the  residue  in  the  filter  add  NH4OH.     Which  salt  is  dissolved1? 
Which  one  is  not  dissolved  ?     Test  filtrate  by  adding  few  drops  of  nitric 
acid.     The  equation  is  : 

(NH3)3(AgCl)2+3HN03==2AgCl  +  3NH4N03. 

(c)  What  salt  is  blackened  upon  filter-paper  1 
Its  formula  being  NH2Hg2Cl,  write  equation, 


110  LABORATORY   MANUAL. 

Experiment  111.— Confirmatory  Tests  for  First  Group. 

Take  25  c.  c.  of  each  of  the  solutions  of  these  first-group  metals,  and 
test  each  hy  adding  a  few  drops  of  each  of  the  following  reagents  : 
KOH  or  NaOH  in  excess. 
NH4OH,  H2S,  H3S04,  KBr,  XI,  and  K2Cr04. 
Write  equation  in  each  case,  and  underline  precipitate ;  as, 
\  AgN03+HCl  =  AgCl  +  HN03. 

Compare  precipitates  when  formed,  as  to  color,  solubility  in  HN03 
and  NH4OH,  and  tabulate  your  results. 

Experiment  112.— Reduction  of  Silver, 

Make  a  solution  of  common  salt  (NaCl)  in  a  test-tube.  Take  1  c.  c. 
of  AgN03  in  another  test-tube,  and  add  to  it  the  salt  solution  until  all 
the  silver  has  been  precipitated.  AgN03 -f  NaCl  =  AgCl-f?  Filter 
and  transfer  to  a  piece  of  charcoal.  Mix  with  the  precipitate  on  the  char- 
coal as  much  sodium  carbonate,  Na2C03.  Heat  before  the  blow-pipe 
until  you  get  a  bead  of  metallic  silver.  How  can  NaCl  form  AgCl  as 
well  as  HC1  does  ?  How  do  sodium  carbonate  and  charcoal  take  away 
chlorine  from  the  silver  ? 

Test  the  silver  bead  with  H2S04,  HC1,  and  HN03,  to  ascertain  which 
will  dissolve  it  most  readily.  Put  the  bead  in  a  test-tube,  and  add 
a  few  drops  of  acid,  and  heat.  As  soon  as  you  are  satisfied  as  to  whether 
chemical  action  takes  place  or  not,  pour  in  water  and  rinse.  In  the  same 
way  try  the  other  acids.  Which  is  the  best  solvent  of  silver  ? 

Can  you  reduce  lead  or  mercury  in  the  same  way  ? 

Optional  Experiment  113. — Pure  Silver  from  Silver  Coin. 

Put   a  small  silver  coin  into  a  beaker  and  cover  it  with  HNOQ  and 

>.     .      . 
water,  about  equal  parts.     Heat  gently  until  the  coin  is  dissolved.     The 

blue  color  is  due  to  copper.  Why  is  copper  used  in  silver  coins  1  Add 
HC1  until  no  more  silver  can  be  precipitated.  Filter,  and  wash  by  pour- 
ing water  over  it  until  the  liquid  runs  through  colorless.  Transfer  the 
silver  chloride  to  a  porcelain  dish,  add  some  clippings  or  grains  of  zinc,  and 
cover  with  dilute  H2S04.  What  gas  will  be  generated  by  zinc  and  sul- 
phuric acid  1  As  this  gas  escapes  it  finds  an  element,  Cl,  for  which  it  has 
strong  affinity.  It  unites  with  Cl  to  form  HC1,  leaving  the  silver  free. 
In  an  hour  the  silver  will  all  be  set  free,  and  appear  as  a  black,  spongy 
mass.  It  may  be  collected  and  melted  into  a  globule  on  charcoal,  or  it 
may  be  dissolved  in  HN03  to  form  silver  nitrate.  The  common  name 
of  silver  nitrate  is  lunar  caustic.  For  what  is  it  used  1 


112  LABORATORY   MANUAL. 

Experiment  114. — Silver-printing. 

Float  a  small  sheet  of  paper  upon  a  solution  of  silver  nitrate  in  a  dark 
room.  Dry  in  the  dark,  and  then  lay  a  fern  leaf,  a  piece  of  lace,  or  a 
photographic  negative  upon  the  silver  paper. 

Cover  with  glass  and  place  in  the  sunshine  until  the  exposed  parts 
are  blackened.  Kinse  with  water  in  the  dark  room,  and  then  for  a  few 
minutes  in  a  solution  of  hyposulphite  of  sodium,  to  dissolve  out  all  the 
silver  not  acted  upon  by  light. 

Then  rinse  for  some  time  in  water  and  dry. 

How  long  did  it  take  to  blacken  the  silver  salt  in  the  sunlight  ? 

Could  photographic  negatives  be  made  by  use  of  this  salt? 

If  not  practicable,  what  salt  is  used  because  so  quick  that  it  is 
instantaneous  1 

Is  silver  nitrate  used  at  all  in  photography  ? 

Experiment  115.— Silver-plating. 

(For  the  Teacher.) — Let  the  teacher  prepare  cyanide  of  silver  and 
plate  a  key  or  some  such  article  to  illustrate  the  process  of  silver-plating. 

See  Gore's  "  Electro-Metallurgy  "  for  process,  solutions,  etc. 

What  force  separates  the  silver  from  cyanogen  1 

Upon  which  pole  of  the  battery  does  silver  collect  ? 

Why  is  it  best  to  keep  a  piece  of  metallic  silver  upon  the  other 
pole? 

How  is  this  silver  affected  ? 

Why  does  the  silver  not  adhere  where  there  is  grease  or  finger-marks 
on  the  article  to  be  plated  ? 

Write  a  full  account  of  the  process. 

Experiment  116.— Reduction  of  Lead  by  Zinc. 

Suspend  a  piece  of  metallic  zinc  in  a  solution  of  lead  nitrate  or 
acetate.  Allow  it  to  stand  over  night.  Zinc  will  exchange  with  the  lead 
of  the  solution.  Notice  the  tree-like  growth  of  gray  deposit  of  lead. 
Test  the  solution  for  zinc  as  follows  : 

First,  to  precipitate  the  lead  not  yet  exchanged,  add  alcohol,  and  then 
sulphuric  acid.  Alcohol  prevents  the  sulphate  of  lead  from  being  dis- 
solved in  water  as  fast  as  precipitated  by  the  acid.  Now  filter  out  the 
sulphate  of  lead  and  neutralize  the  filtrate  with  ammonia  and  then  add 
(NH4)2S,  and  the  zinc  will  be  precipitated  as  a  white  sulphide. 

What  color  is  lead  sulphide  ? 

Complete  the  equation  Zn-f  Pb(N03)2  =  ? 


114  LABORATORY    MANUAL. 

Experiment  117.— Distinction  between  Mercuric  and  Mercurous  Salts. 

Dissolve  a  globule  of  mercury  in  HN03,  using  more  of  the  acid  than 
necessary  to  dissolve  the  globule,  and  you  will  have  mercuric  nitrate, 
Hg(N03)2.  (Remove  gold  rings  from  the  lingers  while  working  with 
mercury.) 

Partly  dissolve  another  globule  in  HN03  and  pour  off  the  acid 
into  another  test-tube,  and  you  will  have  mercurous  nitrate,  Hg.,(N03)2. 
Dilute  these  solutions  with  water  and  apply  the  following  tests  : 

(a)  Try  a  little  of  each  with  H2S.     What  is  the  result? 

(b)  Try  a  little  of  each  with  NaHO. 

(c)  Try  a  little  of  each  with  HC1.      Which  one  gives  a  precipitate  ? 
(cF)  Introduce  a  clean  strip  of  copper  or  a  clean  copper  coin  into  each. 

Rub  with  cloth  after  it  becomes  blackened.  Try  zinc.  Why  is  the  pre- 
caution given  above  in  regard  to  gold  rings  ? 

Sum  up  the  differences  between  mercuric  and  mercurous  compounds. 

It  will  be  noticed  that  mercurous  compounds  belong  to  the  First  Group, 
and  that  mercuric  compounds  belong  to  the  Second  Group.  Also  that 
lead  is  not  all  or  completely  precipitated  by  HC1 ;  so  in  the  key  at  the 
end  of  the  book  it  is  to  be  found  also  in  the  Second  Group. 

Experiment  118.— Reduction  of  Mercury. 

Put  a  small  fragment  of  the  common  ore  of  mercury  HgS,  called  cin- 
nabar, into  a  hard  glass  tube  open  at  both  ends.  Hold  the  tube  inclined 
upward  in  the  flame,  and  heat  strongly.  What  gas  is  given  off?  What  in- 
dication of  metallic  mercury  ?  Why  must  the  tube  be  open  at  both  ends? 

SECOND  GROUP   METALS. 
Sb,  Sn,  Hg",  Bi,  Cu,  and  Pb. 

NOTE. — In  testing  solutions  of  unknown  substances,  the  lead,  if  present,  is  not 
all  precipitated  by  HC1,  and  hence  must  be  looked  for  again  in  the  Second  Group. 

Experiment  119.— Separation  of  Second  Group  into  Two  Divisions. 

Have  solutions  of  the  following  reagents  ready  for  use :  SbCl3  (HC1 
solution),  SnCl2,  Hg(N03)2,  Pb(N03)2,  Bi(N03)2,.and  Cu  (N03)2.  Add 
dilute  HC1  to  all  but  PbN'03  until  each  gives  acid  reaction  with  litmus. 

(a)  To  5  c.  c.  of  SbCl3  add  quantity  strong  solution  H2S  and  filter. 
Wash  precipitate  thoroughly  in  quantity  of  cold  H20,  and  throw  away 
the  nitrate. 

(b)  Pierce  the  filter-paper  with  a  platinum  wire  or  a  glass  rod,  and 
wash  the  precipitate  through  into  an  evaporating  dish. 

(c)  Add  (NH4)2S2,  yellow  ammonium    sulphide,  using  as  little   as 
possible,  and  stir  for  some  time.     The  precipitate  should  dissolve. 


116  LABORATORY    MANUAL. 

Experiment  120. 

Perform  the  same  experimeirts,  using  SnCl2  instead  of  SbCl3.  Does 
the  precipitate  dissolve  at  point  (c)  1 

Experiment  121. 

Perform  the  same  experiments,  using  Cu(N03)2.  Does  precipitate 
(CuS)  dissolve  in  (NH4)2S2  ? 

Experiment  122. 

Do  the  same  with  Bi(N03)2,  Hg(N03)2,and  Pb(N03)2.  Separate 
metals  of  the  Second  Group  into  two  divisions,  based  on  the  solubility  of 
their  sulphides  in  (NH4)2S2. 

Experiment  123.— Identification  of  Sb  in  Solution  of  its  Salts. 

Add  H2S  to  SbCl3  solution.  Note  the  color  of  the  precipitate  in 
(NH4)2S2  and  re-precipitate  by  using  dilute  HC1. 

NOTE. — The  "spot  "  test  for  antimony  is  one  of  the  best,  and  may  be  per- 
formed by  the  teacher  if  thought  desirable.  Arsenic  also  belongs  to  this  group. 
This  and  many  other  elements  may  be  studied  with  profit  in  more  advanced  classes. 

Experiment  124.— Identification  of  Sn. 

(d)  Wash  thoroughly  on  the  filter  the  precipitate  obtained  by  adding 
H2S  to  SnCl2  solution.  What  is  the  color  of  this  precipitate? 

(//)  Dissolve  in  yellow  ammonium  sulphide  and  re-precipitate  with 
HC1,  as  in  the  case  of  antimony.  What  is  the  color  of  this  precipitate  ? 
Is  it  the  same  as  the  color  obtained  in  (a)  ? 

Experiment  125.— Sn,  Continued. 

(«)  Place  some  tin  salt  upon  charcoal  and  use  the  oxidizing  flame. 
Sn03  is  formed  about  the  edge  of  the  substance.  Note  its  color  when 
hot  and  when  cold.  Moisten  the  oxide  when  cold  with  cobaltous  nitrate, 
and  heat  again.  Note  color. 

(b)  To  solution  of  SnCl2  add  some  HgCl2.     Color  of  precipitate? 
Allow  contents  of  tube  to  stand,  and  note  change.     Do  you  find  any 
metallic  Hg  1 

(c)  Decompose  solution  of  SnCl2  by  means  of  electric  current. 

(d)  Add  NH4OH  to  solution  of  SnCl2  and  put  in  a  strip  of  Zn. 

NOTE. — The  separation  of  Sn  from  Sb  may  be  left  for  a  more  advanced  course. 


118  LABORATORY   MANUAL. 

IDENTIFICATION    OF  -METALS    IN    SECOND    DIVISION    OF   SECOND 

GROUP. 

Cu,  Bi,  Hg",  and  Pb. 

Experiment  126,— Copper, 

(a)  Dissolve  a  small  bit  of  copper  in  nitric  acid.     What  gas  is  given 
off?     Color  of  solution  ?     To  what  salt  is  this  due  ? 

(b)  Heat  a  little  copper  sulphate  in  evaporating  dish.     Note  change 
of  color  as  water  is   driven   off.      Try  to  restore  color  with  a  drop  of 
alcohol.     Try  water. 

(c)  Precipitate  CllS  from  solution  of  Cu(N03)2.     Wash  and  dissolve 
precipitate  on  the  filter-paper  with  hot  HN03. 

Evaporate  the  solution  thus  obtained  under  the  hood  in  the  labora- 
tory or  outside  the  window. 

Dissolve  the  residue  in  small  quantity  of  water. 

If  a  slight  white  precipitate  is  formed  in  the  water  at  this  point, 
add  a  drop  or  two  of  HN03  until  it  disappears. 

Add  NH4OH.     Blue  solution  indicates  copper. 

Experiment  127.— Confirmatory  Tests  for  Copper. 

(a)  Dip  a  borax  bead  into  a  solution  of  some  copper  salt,  and  heat 
before  oxidizing  flame.  The  bead  should  be  green  when  hot,  blue  when 
cold. 

(&)  Place  two  drops  of  copper  solution  on  a  clean  knife-blade.  In 
a  few  moments  dip  the  blade  into  water  and  note  copper,  reduced  from 
its  salt  by  iron. 

(c)  Allow  the  clean  copper  wire  to  stand  for  a  few  minutes  in  solu- 
tion of  AgN03. 

What  is  formed  on  its  surface  ? 
Explain. 

(d)  Place  a  drop  of  mercury  in  a  watch-glass  and  cover  with  concen- 
trated solution  of  AgN03.     Explain  the  phenomenon. 

Experiment  128.— Bismuth. 

Dissolve  the  sulphide  formed  by  adding  H2S  to  the  solution  in 
HN03,  and  add  excess  of  NH4OH. 

White  precipitate  is  Bi(OH)3. 

(For  experiments  on  Hg"  and  Pb,  see  those  metals  under  First  Group;) 

To  a  nearly  neutral  solution  of  a  bismuth  salt  add  a  large  quantity 
of  water,  and  a  precipitate  will  be  formed. 


120  LABORATORY    MANUAL. 

Experiment  129.— Separation  of  Metals  in  Second  Division  of  Second 

Group. 

(a)  Dissolve  precipitate  left  after  removing  Sb  and  Sn  from  a  solu- 
tion containing  salts  of  all  the  Second  Group  metals,  in  hot  nitric  acid. 
In  case  it  will  not  all  dissolve,  the  residue  may  be  mercury  (Hg"),  which 
is  now  in  the  form  of  HgS,  and  may  be  dissolved  in  nitro-hydrochloric 
acid.  To  this  solution  add  SnCl2.  The  reaction  which  occurs  is  the 
test  for  Hg". 

(/>)  Evaporate  solution  (ci)  to  dryness,  and  dissolve  residue  in  water. 
If  a  white  precipitate  forms,  add  a  drop  or  two  of  HN03  to  clear.  To  this 
solution  add  a  drop  or  two  of  H2S04.  A  white  precipitate  shows  that 
Pb  is  present  and  H2S04  should  be  added  until  all  Pb  is  precipitated  as 
PbS04.  To  the  filtrate  at  this  point,  or  in  case  there  is  no  lead,  to 
solution  obtained  in  (/>)  add  excess  of  NH4OH.  White  precipitate 
denotes  Bi.  Blue  colored  solution  indicates  copper. 

Write  the  equation  for  the  reaction  when  bismuth  is  present. 

THIRD   GROUP   METALS. 
Fe,  Cr,  Al,  Ni,  Co,  Mn,  Zn. 
Experiment  130.— Fe. 

Dissolve  a  piece  of  iron  wire,  two  inches  long,  in  HC1  in  a  test-tube. 
Write  the  equation  for  the  reaction. 

NOTE. — Iron  acts  as  a  base  in  forming  two  series  of  salts,  the  ferrous  and  the 
ferric.  The  ferrous  salts  may  be  identified  as  such  by  adding  a  few  drops  of  K3FeCyfi, 
potassium  ferricyanide,  which  gives- a  dark-blue  precipitate,  Fe3(FeCy6)3,  ferrous 
ferricyanide. 

Test  a  portion  of  the  solution  just  made  and  dilute  with  water,  and 
determine  whether  the  salt  contained  is  Fe2Cl6  or  FeCl2. 

Experiment  131,— Fe. 

Dilute  a  portion  of  the  remaining  solution  prepared  in  the  preceding 
experiment,  and  add  two  or  three  drops  of  strong  nitric  acid.  Is  there 
any  change  of  color  1 

Boil  carefully  for  a  few  moments  and  cool.  Test  a  small  portion  of 
the  contents  of  your  test-tube  with  a  few  drops  of  K3FeCy6. 

Do  you  obtain  the  same  results  as  before  ?  This  shows  the  salt  of 
iron  to  be  in  the  ferric  condition. 

The  action  of  HN03  upon  Fe  in  changing  it  from  the  ferrous  (Fe") 
to  the  ferric  (Fe'")  condition  may  be  represented  in  the  following  equa- 
tion :  2Fe  +  8HN03  =  Fe2(N03)6  +  2NO  +  4H80. 


122  LABORATORY    MANUAL. 

Experiment  132.— Fe. 

Dissolve  two  or  three  crystals  of  FeS04  in  water  in  your  test-tube. 

NOTE. — Ferrous  salts,  on  being  exposed  to  the  air,  become  coated  with  a  layer 
in  the  ferric  condition.  Hence,  in  getting  solution  of  FeS04,  after  dissolving  away 
a  portion  from  the  outside,  and  when  the  crystals  have  become  a  clear  green  in 
color,  throw  this  solution  away,  add  fresh  distilled  water,  and  warm  gently  to  hasten 
solution. 

This  solution,  if  used  as  soon  as  prepared,  is  in  a  ferrous  condition,  but  will 
become  ferric  on  standing. 

Test  a  portion  of  the  solution  with  K3FeCy6,  as  before. 

Separate  the  solution  into  several  portions,  and  test  with  K4FeCy6 
and  with  KCyS. 

Xanie  these  substances. 

Oxidize  some  of  the  FeS04  to  Fe2(S04)3,  and  test  with  the  three  re- 
agents named  above. 

Make  a  table  showing  results  obtained  witli  each  in  ferrous  and 
ferric  solutions. 

Experiment  133.— Fe, 

(a)  Make  a  fresh  solution  of  FeCl2  as  in  Experiment   130,  and  test 
a  small  portion  with  K4FeCy6,  K3FeCy6,  and  KCyS. 

(b)  Pass  Cl  gas  into  the  remainder.      [This  should  be  done  under 
the  hood.] 

Test  with  the  above  three  reagents,  and  compare  results  with  those 
tabulated  from  preceding  experiment. 

What  is  the  effect  of  Cl  upon  ferrous  salts  ? 
Complete  the  equation,  2FeCl2  +2C1  =  ? 

Experiment  134. — Fe. 

Under  the  hood  pass  H2S  gas  into  dilute  solution  of  Fe2Cl6,  until 
a  few  drops  of  K3FeCy6  reveals  the  presence  of  the  salt  in  the  ferrous 
condition. 

Complete  the  equation,  Fe2Cl6+H2S  =  2FeCl2+? 

In  the  analysis  of  unknown  solutions,  what  reagent  is  added  to  pre- 
cipitate Second  Group  metals  ? 

Is  some  of  it  left  in  the  nitrate  ? 

How  can  you  tell  1 

If  iron  salts  are  present,  what  is  their  condition,  ferrous  or  ferric, 
when  you  are  ready  to  add  the  Third  Group  reagents  1 

If  you  began  with  a  ferric  salt,  what  has  produced  the  change  1 


124:  LABORATORY    MANUAL. 

Experiment   135. 

Add  NH4C1  and  NH4OH  to  solution  of  ferrous  sulphate.  Note 
result. 

To  solution  of  ferrous  sulphate  add  a  few  drops  of  strong  HN03,  and 
heat.  To  this,  when  cool,  add  NH4C1  and  NH4OH  as  before.  Compare 
results.  Explain. 

Thus  we  see  that  if  iron  is  present  in  an  unknown  solution  in  the 
form  of  ferric  salts,  it  is  reduced  to  the  ferrous  condition  by  addition  of 
the  Second  Group  reagent  H2S,  and  must  be  oxidized  to  the  ferric  con- 
dition by  the  addition  of  HN03  and  application  of  heat  before  it  will  be 
precipitated  as  Fe(OH)6  by  NH4C1  and  NH4OH. 

Experiment  136.— AL 

Make  a  solution  of  alum.     Add  NH4C1  solution  and  NH4OH. 
This  precipitates  Al  as  A12(OH)6. 

Note  character  of  the  precipitate  and  compare  with  Fe2(OH)6  formed 
in  the  same  manner.  What  is  alum  ? 

Balance  the  equation  A12C16  +6NH4OH  =  ? 

Experiment   137.— Cr. 

To  a  solution  of  chrome  alum,  K2Cr2(S04)2,  add  some  NH4C1  solu- 
tion and  some  NH4OH.  Cr  is  precipitated  as  Cr2(OH)6.  Compare 
precipitate  with  those  of  iron  and  Al  produced  in  the  same  manner. 

Fe,  Al,  and  Cr  are  thus  seen  to  resemble  each  other  in  this,  that  they 
form  precipitates  (hydrates)  upon  the  addition  of  NH4C1  and  NH4OH 
to  solutions  of  their  salts. 

NOTE  a. — As  has  been  shown,  iron  must  be  in  the  ferric  condition. 

NOTE  b. — Cr  will  not  be  precipitated  as  a  hydrate  at  this  point,  unless  it  is 
present  as  a  base.  To  insure  this,  HC1  and  H2S  must  be  added  to  solutions  which 
we  are  analyzing,  even  if  there  are  no  First  or  Second  Group  metals  present. 

NOTE  c. — Cr  may  be  reduced  to  the  condition  of  a  base  in  still  another  man- 
ner, if  preferred.  In  case  solution  with  which  you  begin  is  K^Cr.,  04,  add  a  few  drops 
of  HC1  and  a  few  drops  of  C2H60,  boil  and  cool.  This  gives  Cr2Cl6. 

1.  Hence,  in  examining  solutions   for  Fe,  Cr,  and  Al,  we  add  HC1 
and  H3S  to  our  solution  to  remove  First  and  Second  Group  metals. 

2.  To  the  nitrate,  if  any  are  present,  or  to  this  solution,  we  add  a  few 
drops  of  strong  HN03,  boil  until  no  odor  of  H2S  is  given  off,  and  cool. 

3.  When  cold,  we  are  ready  for  the  Third  Group  reagents,  which  are 
NH4C1,  NH4OH,  and  (ffH4)2S. 

4.  As  has  been  seen,  the  addition  of  the  first  two  serves  to  separate 
Fe,  Al,  and  Cr  from  the  remainder  of  the  Third  Group,  Ni,  Co,  and  Mn, 
which  will  be  precipitated  as  sulphides  upon  the  addition  of  (N"H4)2S. 


126  LABORATORY   MANUAL. 

Experiment  138.— To  Separate  Metals  of  the  Third  Group  into  Two 

Divisions. 

Treat  a  solution  containing  salts  of  all  three  metals  of  this  group  as 
suggested  for  iron  and  chromium.  (See  1  and  2  above.) 

Experiment  139. 

Add  NH4C1  and  NH4OH,  and  filter.  Test  nitrate  to  see  that  all  of 
the  Al,  Fe,  and  Cr  is  precipitated.  Filtrate  contains  Ni,  Co,  Mn,  and  Zn. 
To  this  add  (NH4)2S  and  obtain  NiS,  CoS,  MnS,  and  ZnS. 

Experiment  140.— To  Separate  and  Identify  Fe,  Al,  and  Cr. 

Using  precipitate  obtained  on  filter  in  preceding  experiment,  wash 
through  into  glass  beaker  or  large  test-tube,  add  quite  a  quantity  of  KOH, 
and  boil  for  some  time. 

The  KOH  dissolves  the  A12(OH)6.  Set  aside  this  solution  to  test 
for  Al.  The  remaining  precipitate,  if  there  is  any,  consists  of  Fe2(OH)6 
or  Cr2(OH)6,  or  both. 

Experiment  141. 

Separate  this  precipitate  into  two  portions  and  dissolve  one  in  HC1, 
and  employ  test  for  iron  given  in  Experiment  132.  Fuse  the  other 
portion  of  the  precipitate  on  platinum-foil  with  sodium  or  potassium 
nitrate  and  carbonate. 

Dissolve  the  mass  obtained  (Na2Cr04)  in  water,  and  apply  tests  for 
chromium  as  shown  in  the  following  experiment : 

Experiment  142. — Cr  Confirmatory  Tests. 

To  a  portion  of  the  solution  of  Na2Cr04  obtained  in  the  preceding 
experiment,  add  acetic  acid  until  the  solution  gives  an  acid  reaction 
with  litmus. 

To  this  add  lead  acetate  solution. 

Note  color  of  precipitate  (PbCr04). 

"Write  equation. 

Experiment  143. 

Borax  bead  is  given  an  emerald-green  color. 

Experiment  144.— AL 

Acidulate    the   potassium   solution    of  Al   obtained    in  Experiment 
140  with  HC1.     To  this  add  (NH4)2C03. 
Precipitate  A12(OH)6  indicates  Al. 


128  LABORATORY   MANUAL. 

Experiment  145. — To  Separate  and  Identify  Ni,  Co,  Mn,  and  Zn. 

After  removing  metals  of  the  first  division  of  the  Third  Group,  add 
(NH4 )  2  S,  and  sulphides  of  Ni,  Co,  Mn,  and  Zn  are  precipitated. 

Filter,  and  save  nitrate  to  examine  for  Fourth  Group  metals. 

Treat  precipitate  on  the  filter  with  dilute  HC1. 

The  black  residue  is  Co  or  Ni,  or  both,  and  filtrate  contains  Zn  or 
Mn,  or  both. 

Experiment  146. 

Test  some  of  the  residue  with  borax  bead  and  blow-pipe. 
A  blue  bead  indicates  Co. 

If  Ni  is  present  alone,  the  bead  will  have  a  brown  color  when  hot, 
and  yellow  when  cold,  in  the  oxidizing  flame  only. 

Experiment  147. 

Dissolve  the  residue  in  nitro-hydrochloric  acid. 

If  Ni  is  present  alone,  Ni(OH)2  will  be  precipitated  upon  the  addi- 
tion of  a  few  drops  of  ammonia. 

The  color  of  the  precipitate  is  green. 

Experiment  148. 

If  Co  and  Ni  are  both  present,  the  tests  interfere  with  each  other  to 
some  extent,  and  the  method  of  separating  them  is  omitted. 

Experiment  149. 

The  filtrate  from  Co  and  Ni,  which  may  contain  Mn  or  Zn,  or  both, 
contains  them  in  the  form  of  ZnCl2  and  MnCL. 

Boil  the  filtrate  to  expel  H2S,  then  cool  it,  and  add  a  decided  excess 
of  KOH. 

Allow  to  stand  for  some  time.  If  Mn  is  present,  it  will  be  precipi- 
tated as  Mn(OH)2. 

Experiment  150, 

With  acetic  acid  acidulate  the  filtrate  from  this.  Add  (NH4)2S, 
and  ZnS  is  precipitated. 

ZnS  is  insoluble  in  acetic  acid. 

Note  the  color  of  the  precipitate. 

How  does  it  differ  from  the  other  sulphides  of  this  division  of  the 
group  ? 


130  LABORATORY    MANUAL. 

Experiment  151.— Fourth  Group  Metals. 

Place  a  small  piece  of  freshly  burned  quick-lime  in  a  large  evaporat- 
ing dish,  and  pour  on  water  until  the  lime  is  nearly  covered. 

Note  the  action  and  describe  it. 

After  the  action  ceases,  is  there  as  much  water  present  as  before  1 
What  has  become  of  it  1     What  substance  is  formed  1     What  is  mortar  1 

Experiment  152. 

Try  CaC03,  Ca(OH)2,  and  CaCl2  as  to  their  solubility  in  water. 

In  making  lime-water  for  experiments  with  C02,  what  substance 
was  dissolved?  What  substance  is  formed  when  C02  is  passed  into 
lime-water  1  Is  it  soluble  ?  How  do  you  know  1 

Experiment  153. 

Heat  moderately  some  powdered  gypsum   (CaS042H20)   in  a  hard 


is  test-tube  or  evaporating  dish.  What  happens  ?  See  if  the  pow- 
der which  is  left  will  harden  when  made  into  a  paste  by  the  addition 
of  a  little  water.  Explain  the  action.  What  is  "  plaster  of  Paris  "  ? 

Experiment  154. 

Fasten  a  small  quantity  of  asbestos  in  a  loop  of  platinum  wire.  Dip 
into  a  solution  of  CaCl2  and  heat  in  the  non-luminous  Bunsen  name  in 
front  of  the  spectroscope.  Make  a  drawing  of  the  spectrum  in  your  note- 
book. Is  the  spectrum  continuous  or  one  of  the  "  bright-line  "  variety  ? 

Compare  also  spectrums  of  SrCl2  and  BaCl2.  What  color  is  im- 
parted to  the  flame  in  each  instance  ? 

Experiment  155. 

Make  a  solution  of  CaCl2  by  dissolving  a  small  piece  of  marble 
(CaC03)  in  HC1. 

Write  the  equation. 

Add  NH4OH,  NH4C1,  and  (NH4)2C03. 

The  substance  formed  is  a  carbonate. 

Write  the  equation. 

Thus  we  see  that  Ca  is  precipitated  as  CaC03  in  the  presence  of 
NH4C1  and  NH4OH.  The  chlorides,  hydroxides,  and  sulphides  of  Ca  are 
not  precipitated  when  HC1,  NH4OH,  or  (NH4)2S  are  added  to  precipitate 
metals  of  the  first  three  groups.  Hence  this  serves  to  form  the  basis  of 
a  new  group,  the  fourth. 

Ba  and  Sr  are  precipitated  in  the  same  way. 


132  LABORATORY   MANUAL. 

Experiment  156. 

To  the  filtrate  from  the  Third  Group  (Experiment  145)  or  to  a  solution 
containing  no  First,  Second,  or  Third  Group  metals,  add  NH4OH,  NH4C1, 
and  (NH4)2C03. 

Precipitate  indicates  Ba,  Ca,  Sr,  or  any  two  of  them,  or  all  (see  Ex- 
periment 147).  They  are  precipitated  as  carbonates. 

Write  the  equations,  using  the  solutions  of  BaCl2,  SrCl2,  and 
CaCl2. 

Save  the  precipitate  on  the  filter  for  use  in  Experiment  158.      Save 
filtrate  for  Experiment  160. 

Experiment  157. 

(a)  Add  some  of  each  of  the  Fourth  Group  reagents  to  solution  of 
BaCl2. 

Filter,  wash  precipitate,  and  try  to  dissolve  it  in  dilute  acetic  acid. 

(b)  Try  same  with  solution  of  CaCl2. 

(c)  Try  same  with  solution  of  SrCl2. 

Experiment  158. 

(a)  Dissolve  precipitate  on  filter  in  Experiment  156.     Substances  in 
solution  at  this  point  are  acetates  of  one  or  more  of  the  Fourth  Group 
metals. 

Write  the  equations. 

(b)  To  a  small  portion  of  the  solution  add  some  K2Cr207  solution. 
Precipitate  denotes  barium.     Yellow  precipitate  is  BaCr04. 

If  barium  is  present,  add  K2Cr207  to  the  rest  of  the  solution,  and 
remove  it  all.  Filter  and  test  filtrate  for  Ca  and  Sr  as  follows : 

(c)  Precipitate  by  (NH4)2C03  and  NH4OH.      Filter  and  wash  pre- 
cipitate. 

Dissolve  it  in  HC1.  Evaporate  (under  hood)  nearly  to  dryness,  and 
make  concentrated  solution  by  adding  a  very  little  water.  Divide 
solution  into  two  portions. 

1.  Test  one  for  Sr  by  adding  solution  of  calcium  sulphate  and  boil- 
ing.     Set  aside  to  cool,  and  in  about  fifteen  minutes  a  precipitate  de- 
notes the  presence  of  Sr. 

2.  Test  the  second  portion  by  adding  a  solution  of  K2S04  and  filter- 
ing, to  make  certain   that  there  is  no  Sr  present.     To  the  nitrate  add 
NH4OH  and  (NH4)2C204  and  precipitate  denotes  Ca. 

Metals  whose  chlorides,  sulphides,  and  carbonates  are  soluble,  and 
hence  not  precipitated  by  any  of  the  group  reagents  used  so  far,  com- 
prise the  Fifth  Group. 


134  LABORATORY    MANUAL. 

THE   FIFTH   GROUP. 
Mg,  K,  Na,  and  NH3. 

Experiment  159.— Magnesium. 

To  the  filtrate  from  Group  Four  (Experiment  156),  or  an  original  so- 
lution, add  Na2HP04.  Precipitate  indicates  Mg. 

It  is  better  to  add  a  little  NH4OH,  and  then  NH4C1,  before  adding 
Na3HP04,  when  a  fresh  solution  of  MgCl2  is  being  tried  for  the  test. 

Experiment  160.— Ammonia. 

Add  KOH  or  NaOH  solutions  to  your  solution,  and  warm  gently. 
Odor  of  NH3  is  a  test.  Note  effect  of  vapor  of  NH3  on  moistened  lit- 
mus paper.  Hold  glass  rod  moistened  with  HC1  near  the  tube  contain- 
ing the  solution. 

What  is  formed  on  the  rod  ? 

Write  the  equation. 

Experiment  161.— Sodium. 

Sodium  is  detected  in  solution  of  its  salts  by  the  color  which  they 
give  to  the  (a)  non-luminous  Bunsen  flame,  and  by  the  spectrum.  Note 
these  carefully.  See  experiments  on  Ba,  Ca,  and  Sr,  and  compare 
spectra. 

(b)  Review  your  experiments  in  making  common  salt  and  in  decom- 
posing water  by  the  action  of  metallic  potassium.  Will  metallic  Na  do 
as  well  1 

Experiment  162.— Potassium,  K. 

Examine  color  given  to  flame,  as  in  Experiment  151,  when  KC1 
solution  is  used.  Examine  its  spectrum. 

These  are  tests  of  K  in  solution  of  its  salts. 

Experiment  163. 

Examine  the  metallic  K  and  Na  kept  in  the  laboratory.  Why  are 
they  kept  under  oil  ?  What  element  must  be  kept  under  water  1 

Experiment  164. 

Place  some  wood  ashes  on  a  filter  and  pour  on  water.     As  fast  as  it 
runs  through  pour  it  back,  over  and  over  again. 
Test  this  solution  for  K. 

Where  does  the  K  in  the  ashes  come  from  1     Give  its  history. 
Add  a  few  drops  of  HC1  to  the  solution.     What  gas  is  given  off? 
What  salt  of  K  is  in  the  solution  ? 


136  LABORATORY   MANUAL. 

WEIGHTS  AND  MEASURES. 
APOTHECARIES'   MEASURE 

60  minims  make  1  fluid  drachm. 

8  fluid  drachms       "  1  fluid  ounce. 

16  fluid  ounces         "  1  pint. 

8  pints  "  1  gallon. 

APOTHECARIES'  WEIGHT. 

20  grains  make  1  scruple. 

3  scruples  "  1  drachm. 

8  drachms  "  1  ounce. 

12  ounces  1  pound. 

COMPARATIVE,   FLUID   MEASURE. 


1 

4 
5 

8 

c. 
c. 
c. 
c. 

c. 
c. 
c. 
c. 

=       17  minims  (nearly). 
=       63       "        =1  drachm 
=      85       "       =  1 
=    136       "       =2       " 

and 
a 

it 

8 
25 
16 

minims. 

c; 
(( 

10 

c. 

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—  _  2 

a 

n 

50 

" 

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c. 

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=  1 

ounce 

cc 

0 

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30 

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50 

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a 

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6 

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=  2 

« 

u 

1 

u 

0 

it 

100 

c. 

c. 

=  1700       " 

=  3 

ti 

tt 

4 

cc 

20 

f'. 

COMPARATIVE,   SOLID   MEASURE. 

1  gram  =  15|  grains. 

5  "  =  77  "  =1  drachm  and  17    grains 

8  "  =  123-J-  "  =2  "  "       3J     " 

10  "  =  154  "  =    2  "    34  " 

20  "  =  308  "  ==    5  "  "8  " 

50  "  =  770  "  =  12  "  "    50  " 

100  "  =  1540  "  =  25  "  "    40  " 

METRIC   SYSTEM. 
LINEAR     MEASURE. 

10  millimetres  (mm.)  =  1  centimetre  (cm.  or  c.)  =       .39  in. 
10  centimetres  =  1  decimetre  (dcm.)  =    3.03   l" 

10  decimetres  =  1  metre  (in.)  =  39.37  " 

10  metres  =  1  dekametre. 

10  dekametrea  =  1  hektometre. 

10  hektometres  —  1  kilometre. 


LABORATORY   MANUAL. 


137 


CAPACITY. 

10  millilitres    =  1  centilitre. 


10  centilitres 
10  decilitres 
10  litres 
10  dekalitres 


1  decilitre. 

1  litre  =  61  cu.  in. 

1  dekalitre. 

1  hektolitre. 


10  hektolitres  ==1  kilolitre. 


WEIGHT. 

10  milligrams  (nig.)  ==  1  centigram  (cgm.). 


10  centigrams 
10  decigrams 
10  grams 
10  dekagrams 
10  hektograms 


=  1  decigram  (dcg.). 

=  1  gram  =  15.43. 

=  1  dekagram. 

=  1  hektogram. 

=  1  kilogram  (kgm.)  = 

2)11* 


Do  not  try  to  translate  these  weights  and 
measures  into  common  or  U.  S.  terms.  One 
can  very  quickly  learn  the  three  units,  the 
Gram,  the  Litre,  and  the  Metre,  and  then  the 
important  divisions  and  multiples  of  these  units. 
Examine  and  use  the  weights  and  measures 
until  they  become  as  familiar  as  ounces,  pints, 
or  pounds.  In  writing  the  fractions  of  a  unit 
use  the  decimal  system,  thus  : 

Millimetre,   0.001 
Centimetre,    .01 
Decimetre,      .1 
Metre,  1. 

A  cubic  centimetre  of  distilled  water  at  4°  C. 
weighs  one  gram.  One  litre  =  a  cubic  deci- 
metre, or  a  thousand  cubic  centimetres  =  1.05672 
U.  S.  quarts. 


H  Z~ 

U 


THERMOMETRIC   RULES. 

To  change  Fahrenheit  degrees  to  Centigrade, 
subtract  32°  and  multiply  the  remainder  by  |.     C.  =  (F.  —  32)  -|. 

To   change  Centigrade   degrees    to   Fahrenheit,  multiply  by  -|  and 
add  32°.     F.  =|C/+  32°. 


138  LABORATORY   MANUAL. 

STANDARD   SOLUTIONS. 

Iodine   Solution  for   Starch   Test. 

To  1  gram  of  metallic  iodine  and  1.5  grams  of  dry  potassium  iodide 
add  100  c.  c.  of  distilled  water.  Stir  with  a  glass  rod  until  all  the  iodine 
dissolves.  The  object  of  the  potassium  iodide  is  to  increase  the  solubility 
of  the  iodine.  This  solution  will  not  keep  well,  and  so  should  be  freshly 
made  for  the  tests  upon  digestion  of  starch. 

Copper  Sulphate    Solution  for  Detection    of    Grape  Sugar — Fehlin(fs 

Fluid. 

(a)  Dissolve  3.46  grams  of  crystallized  cupric  sulphate  in  16  c.  c.  of 
distilled  water. 

(b)  Dissolve  17.3  grams  of  pure  crystallized  potassium-sodium  tar- 
trate  in  60  c.  c.  of  sodium  hydrate,  specific  gravity  1.12. 

Add  (a)  to  (b),  stirring  well ;  then  dilute  to  500  c.  c.  with  distilled 
water.  This  fluid  will  not  keep  long,  so  it  is  best  to  prepare  it  each 
time  it  is  required.  If  kept  cool  and  dark  in  full,  glass-stoppered  bottles, 
it  will  keep  for  a  week  or  more.  It  is  well  to  test  it  by  boiling,  to  see 
if  it  will  reduce  alone,  before  using  it  for  sugar  test.  .005  gram  of  grape- 
sugar  will  completely  reduce  1  c.  c.  of  this  solution. 

REAGENTS. 

Strength  of  Solutions  Usually  Employed. 
Acids : 

Acetic  (vinegar),  30%,  specific  gravity  1.04. 
Hydrochloric  (Muriatic),  24%,  specific  gravity  1.12. 
Hydrosulphuric,  gas  or  saturated  solution. 
Nitric,  32%,  specific  gravity  1.2. 

Sulphuric  (Oil  of  Vitriol),  concentrated  1.8.     Dil.  1  to  5  water. 
Aqua  regia,  1  part  concentrated  HN03  to  4  parts  HC1. 
Ammonium  : 

Chloride,  1  part  crys.  salt  to  8  water. 

"          Carbonate,  1  part  crys.  salt  to  4  water -f  1  part  NH4HO. 
"          Hydrate,  12%,  specific  gravity  .95. 
"          Sulphide,  NH4HO  saturated  with  H2S. 
Barium  Chloride,  1  part  to  10  of  water. 
Hydrate,  1  part  to  20  in  H30. 
Calcium        "        saturated  solution  in  H20. 
"       Sulphate,          u  "  " 


LABORATORY  MANUAL. 

Calcium  Chloride,  1  part  to  8  H20. 

Cobalt  Nitrate,  1  "  "  8  " 

Copper  Sulphate,  1  "  "  8  " 

Ferrous  Sulphate,  1  "  "  5  " 

Ferric  Chloride,  1  "  "  15  " 

Lead  Acetate,  1  "  "  10  " 

Magnesium  Sulphate,  1  "  "  10  " 

Mercuric  Chloride,  1  "  "  16  " 

Potassium  Bichromate,  1  "  "  10  " 

Cyanide,  1  "  "  4  " 

Ferricyanide,  1  "  "  12  " 

Ferrocyanide,  1  "  "  12  " 

Iodide,  1  "  "  20  " 

Sulphate,  1  "  "  12  " 

Sodium  Carbonate,  1  "  "  5  " 

«        Hydrate,  1  "  "  20  " 

«        Phosphate,  1  "  "  10  " 

Silver  Nitrate,  1  "  "  20  « 

Stannous  Chloride,  1  "  "     6  "     +  HC1  (acidulate). 


140 


LABORATORY   MANUAL. 


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144  LABORATORY    MANUAL. 

QUESTIONS   SUGGESTING   FURTHER  STUDY   OF   THE   METALS. 

Of  what  use  is  mercury  in  gold  mining  1  Dissolve  gold-foil  in  a 
globule  of  mercury.  What  is  an  amalgam  ?  What  per  cent,  of  gold  in 
gold  coins?  FeS04  will  precipitate  gold  from  its  solutions,  and  aqua 
regia  will  dissolve  it,  forming  a  chloride.  3FeS04  -f  AuCl3  —  Fe2(S04)3 
+  FeCl3+Au  and  Au  +  3Cl—  AuCl3.  What  is  brass?  How  could 
you  separate  it  into  its  constituents  1  What  is  plumbers'  solder  1  What 
is  the  probability  that  any  of  the  metallic  elements  may  be  found  to  be 
compounds  ?  (See  Professor  William  Crookes'  article  in  forum  for 
May,  1891.)  Why  are  bicarbonate  of  sodium  and  potassium  useful  in 
cake-making  ?  Why  is  yeast  used  in  making  bread  ?  Could  bread  be 
made  light  without  any  of  these  ?  How  ?  How  does  plaster  of  Paris 
differ  from  gypsum  1  How  is  quick-lime  obtained  1  What  metals  are 
useful  in  medicine  ?  If  potassium,  sodium,  and  ammonia  are  so  similar 
in  all  their  uses,  is  there  any  probability  of  the  resolution  of  the  two 
former  into  other  elements,  as  ammonia  is  ?  Has  any  prominent  scientist 
ever  announced  a  belief  in  all  elements  being  some  time  reduced  to  a 
form  of  hydrogen  1  What  metals  are  found  free  in  nature  ?  How  do 
other  metals,  when  in  an  elementary  state,  even  the  common  ones,  com- 
pare in  value  with  gold  ?  Why  is  gold  called  a  precious  metal  1  What 
metals  were  first  known  to  the  ancients  ?  What  uses  could  the  metal 

aluminum  be  put  to  when  cheap  enough  for  general  use  ? 

* 

To  THE  TEACHER. — The  pupil  should  now  be  given  a  number  of  substances  to 
determine  first  the  base,  then  the  acid.  These  should  be  combined  and  the  name 
and  symbol  of  the  salt  written  out.  It  will  be  noticed  that  a  different  method,  to 
some  extent,  has  been  used  in  forming  a  Key,  to  that  used  in  the  experiments  upon 
metals.  The  Key  is  only  a  brief  method,  and  offers  some  advantages  in  rapid  work, 
but  the  pupil  should  confirm  the  results  obtained  here  by  referring  to  the  experi- 
ments in  the  text  for  further  proof  of  his  work, 


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